AN ELEMENTARY 



TEXT-BOOK OF CHEMISTRY. 



WILLIAM GL MIXTER, 

Professor of Chemistry in the Sheffield Scientific School of Tale University. 



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NEW YORK : 

JOHN WILEY & SONS, 

15 Astor Place. 

1889. 



Copyright, 1889, 

BY 

William G. Mixter. 






PREFACE. 



This work is designed for use in schools and colleges. The 
aim is to present the elements of chemistry logically as far 
as possible, so that the student may grasp the fundamental 
principles of the science, and at the same time learn some- 
thing of the chemistry of common things. 

With this aim in view the Periodic Classification has been 
adopted. The acidic and basic groups are treated alternately 
in order to discuss bases and salts early in the course, as well 
as to give constant variety to the character of the experiments 
performed. Compounds of the rare elements are described to 
make evident the reasons for the classification, and also to 
serve as a basis for the summaries of the groups. The de- 
scriptions of the rare elements and their compounds, and of 
the less important compounds of the common elements, are 
in small type. 

Graphic and constitutional formulas are much used, since 
they better represent the properties of compounds than em- 
pirical formulas, and also because they facilitate the study 
of inorganic as well as organic chemistry. The reasons for a 
number of constitutional formulas are given, and in case of 
compounds whose constitution is not understood, cai'e is 
generally taken to state that the constitutional formulas em- 
ployed are assumed from analogy. While graphic formulas 
are very constantly used for common substances for the sake 
of rendering such formulas familiar, it is not intended to 
imply that the use of empirical formulas should be avoided. 



IV PREFACE. 

What to teach beginners in chemistry, and how best to do 
it, are difficult questions. Most students are interested in 
the subject, and study faithfully and enthusiastically, needing 
only guidance and instruction. It is the duty of the teacher 
to help them to think for themselves, to aid in drawing con- 
clusions from experimental data, and to make sure that 
they understand the reasons for the fundamental theories. 
Experience shows that the student is apt to retain a 
general knowledge of the philosophy of chemistry far better 
than a specific knowledge of substances; nevertheless, knowl- 
edge of many compounds is necessary to the discussion and 
comprehension of theories. The work in the laboratory yields 
this knowledge most readily, and at the same time trains the 
powers of observation. 

Numerical data, though essential in descriptions of sub- 
stances, should not as a rule be required for recitation. A 
table of atomic weights, hung in the class-room, will remove 
the temptation to memorize atomic weights. Such a table, 
which can be best exhibited in the form of the periodic 
classification, will early render familiar the groups of the 
elements, and be of service in many other ways. It is 
recommended,, in assigning a lesson, to point out the 
portions requiring special attention and also those parts 
of the text which may be omitted. The difference in 
type used will aid the teacher in his selections, but some 
of the text in large type may properly be regarded as 
matter for reading rather than for recitation. The instructor 
should emphasize the fact that the members of a group form 
analogous compounds with other elements. 

It is preferable to have the work in the laboratory, as far 
as practicable, precede the recitation on a topic. The stu- 
dents can then more readily understand the statements con- 
tained in the text. Questions on the laboratory work and a 
discussion of the results of experiments are perhaps of more 
value than questions on the text. At the beginning of the 



PREFACE. V 

course all experiments should be made before the class, and 
detailed directions given for those to be made by the students. 
Afterwards it will be better to have them depend more upon 
the results derived from their own experiments. Constant 
supervision of laboratory work and aid in manipulation are 
necessary, but the student should be left to obtain results, the 
suggestions coming from the teacher rather in the form of 
questions. Eeference should be made during the laboratory 
work to the text-book for information about substances and 
reactions. 

The author acknowledges with pleasure the valuable aid 
his colleagues have generously given him in the preparation 
of this book. Professor Wm. H. Brewer, who has for many 
years given instruction in the laboratory to beginners, has 
made many suggestions about experiments and methods of 
teaching. Professor Charles S. Hastings has supplied the 
part on Physics of Chemistry and Spectral Analysis, and 
given assistance on portions of the proof. Professor E. H. 
Chittenden read most of the manuscript and made many 
suggestions. Professor H. L. Wells read the part on iron 
and steel. 

Sheffield Laboratory of Yale University, 
New Haven, December, 1888. 



CONTENTS. 



PHYSIOS OF CHEMISTRY. 

Fundamental conceptions, 1. Forms of matter, 5. Balance, 
7. Determination of densities, 9. Crystallography, 12. Mo- 
lecular structure of matter, 22. Temperature and heat, 24. 
Specific heat, 28. Pressure and volume of gases, 31. Laws 
of gases, 33. Kinetic theory of gases, 34. Determination of 
gas densities, 41. 



CHEMISTRY. 

Elements, Atoms, Classification, 46 

Elements, 46. Atomic weight, 47. Symbols, formulas, 47. 
Table of atomic weights, 48. Classification of the elements, 
49. Hydrogen, 50. 

Seventh Group, 56 

Chlorine, 56. Molecules, 61. Molecular weights of com- 
pounds, 63. Bromine, 70. Iodine, 71. Fluorine, 75. Sum- 
mary of the halogens, 77. Manganese, 78. 

Valence, 80 

First Group, 88 

Alkali metals, 83. Atomic volume, 84. Lithium, 84. So- 
dium, 85. Potassium, 87. Rubidium, 89. Ciesium, 89. 

Spectral Analysis, 90 



Vlll CONTENTS. 

PAGE 

First Group — {Continued), 94 

Copper, 94. Silver, 97. Gold, 101. 

Sixth Group, 107 

Oxygen, 107. Ozone, 110. Water, 113. Hydrogen diox- 
ide, 119. Oxides and hydroxides of seventh group, 123. 



Valence of the elements of the seventh group, and constitu- 
tion of their oxygen compounds, 131. Summary of seventh 
group, 135. Oxides and hydroxides of first group, 137. 

Sixth Group— ( Continued), 144 

Sulphur, 144. Selenium, 147. Tellurium, 147. Sulphides 
and hydrosulphides of first group, 152. Oxides and hydroxides 
of sulphur, selenium, and tellurium, 154. 

Bases, Acids, and Salts, 162 

Sixth Group — {Continued) 166 

Constitution of sulphuric acid, 167. Oxides and hydrox- 
ides of sulphur, selenium, and tellurium continued, 169. Sul- 
phates of first group, 172. Chromium, 177. Molybdenum, 
185. Tungsten, 187. Uranium, 190. Summary of sixth 
group, 192. 

Second Group, 195 

Beryllium, 195. Alkali-earth metals, 196. Calcium, 197. 
Strontium, 202. Barium, 203. Magnesium, 206. Zinc, 209. 
Cadmium, 213. Mercuiy, 214. 

Fifth Group, 222 

Nitrogen, 222. Oxides and hydroxides, 230. Nitrates, 236. 
Constitution of oxygen compounds of nitrogen, 248. The 
Atmosphere, 249. 

Phosphorus, 252. Halides, 258. Oxides and hydroxides, 
260. Phosphates, 265. 



CONTENTS. IX 

PAGE 

Arsenic, 268. Antimony, 276. Bismuth, 281. Vanadium, 
283. Niobium, 285. Didymium, 286. Samarium, 287. Tan- 
talum, 288. Summary of fifth group, 288. 

Third Group, . . . • . . - 291 

Boron, 291. Aluminum, 296. Gallium, 303. Indium, 304. 
Thallium, 305. Scandium, 308. Yttrium, 309. Lanthanum, 
309. Erbium, 310. Ytterbium, 311. Summary of third 
group, 311. 

-Fourth Group, 313 

Carbon, 313. Compounds of, 321. Isomerism, polymerism, 
322. Carbonic acid and carbonates, 327. Cyanogen com- 
pounds, 341. Methane and derivatives, 346. Constitution of 
derivatives of methane, 349. Compound ammonias, 352. De- 
rivatives of ethane, 354. 

Silicon, 363. Glass, 369. Porcelain, brick, hydraulic 
cement, 371. Titanium, 372. Zirconium, 374. Cerium, 375. 
Thorium, 376. Germanium, 377. Tin, 378. Stannous com- 
pounds, 380. Stannic compounds, 381. Lead, 385. Sum- 
mary of fourth group, 392. 

Eighth Group, 394 

Iron, 394. Varieties, 395. Manufacture, 396. Properties of 
iron and steel, 399. Ferrous compounds, 400. Ferric com- 
pounds, 405. Iron cyanides, 408. 

Cobalt, 414. Nickel, 417. Ruthenium, 421. Rhodium, 422. 
Palladium, 423. Osmium, 425. Iridium, 426. Platinum, 427. 
Summary of eighth group, 431. 

Atomic Theory, 435 

Dalton hypothesis, 435. Law of multiple proportions, 436. 
Law of definite proportions, 437. Determination of atomic 
weights, 438. Law of Dulong and Petit, 440. Data required 
for an atomic weight, 441. 

Periodic Law, 443 



PHYSICS OF CHEMISTRY. 



Fundamental Conceptions. 

Physical Magnitudes. — There are three fundamental con- 
ceptions of quantity which furnish the means for measuring 
all magnitudes involved in physical science, namely, of space, 
of time, and of matter. Our ideas, whether innate or derived 
from experience, are sufficiently precise concerning the first 
two quantities to enable us to presume a sufficient knowledge 
of their natures for all practical purposes; the last, however, 
requires more particular consideration. 

Matter. — The existence of matter, and of matter alone, is 
demonstrated to us by the direct evidence of our senses. We 
infer the existence of forces, of electricity, etc., by observa- 
tion of changes going on in matter, but we have no sense- 
organs which betray the existence of any of them to the mind 
directly. For the discussion of purely chemical phenomena 
we may define matter as anything which can be weighed, and 
the quantity of matter as proportional to its weight. 

The one essential property of matter from the chemical 
standpoint is its indestructibility. From this property, which 
forms the logical foundation of the science of chemistry, we 
may deduce the practical rule: 

In any closed space the total quantity of matter is i. 'trad- 
able, irrespective of the clia ayes of form which it may undergo. 

For example, if we have a closed vessel containing a quan- 
tity of ice, we may by heating it convert it firsl into water, 
1 



2 PHYSICS OF CHEMISTRY. 

then into steam, then into its component gases, hydrogen and 
oxygen, and then, probably, into simpler forms of these ele- 
ments; but with all these changes the total quantity of matter 
remains unchanged. 

Fundamental Units. — The unit of length now almost uni- 
versally employed in science is the centimeter (equals 10 milli- 
meters), and is the unit which we adopt. 

The unit of time is the second. 

The unit quantity of matter, or unit of mass, is a quantity 
of matter equivalent to that contained in a cubic centimeter 
of water at its maximum density. It is called a gram. 

Derived Units. — As has already been noted in the first para- 
graph, all quantities measured in physical science can be ex- 
pressed in terms of these three fundamental units. For ex- 
ample, an area is expressed in terms of the unit of length 
squared, that is, in square centimeters: a velocity, in terms of 
the unit of length divided by the unit of time; as, for example, 
a velocity of ten, means a change of place at the rate of ten 
centimeters per second, etc. These are examples of derived 
units. 

Among the most important derived units employed in 
chemistry are: 

1° Doisity. — This is defined as the quantity of matter di- 
vided by the space it occupies. Since the definition of unit 
quantity of matter tells us that in a body of water the number 
of units of mass is equal to the number of cubic centimeters 
it occupies, the density of water is unity. As experiment 
shows that a given bulk of mercury contains 13.6 times as 
much matter as an equal bulk of water, the density of mer- 
cury is 13.6. 

It is obvious that these values depend upon the arbitrary 
condition that a unit volume of water shall contain a unit of 
matter. In the English system of units, if we take the cubic 



FUNDAMENTAL CONCEPTIONS. 3 

foot as the unit of volume and the pound as the unit of mass, 
the density of water (since a cubic foot contains 62.5 pounds 
of matter) is 62.5. In any system, however, the ratios of the 
densities are the same; hence, as the density is very character- 
istic of liquids and solids, these ratios taken with respect to 
water are made to replace the inconvenient densities in sys- 
tems other than the metric. Such ratios are called specific 
gravities : they are obviously numerically equal to the densi- 
ties in the metric system. As the last-named system is 
adopted in this book, the term has no place in it. 

2° Motion. — A quantity of motion is equal to the product 
of the mass of a moving body by its velocity. Thus, the mo- 
tion of a kilogram of matter moving one centimeter per second 
is equal to the quantity of motion of a gram moving a thou- 
sand centimeters per second, which is equal to a thousand 
units of motion. 

3° Force. — The cause of change of motion is called force, 
and its measure is the rate at which it can produce the change. 
A unit force^ is, therefore, a force which can increase the ve- 
locity of motion of a gram one centimeter per second, or the 
velocity of motion of a kilogram one one-thousandth of a centi- 
meter per second. It is called a dyne. If a gram of matter 
be allowed to fall freely, it is found that its velocity increases 
at the rate of 980 centimeters per second, hence the weight of 
a gram is 980 dynes. 

4° Pressure. — A force may be applied at such a small por- 
tion of the surface of a body that this portion may conven- 
iently be regarded as a point; for example, when a moving 
billiard-ball strikes another, the surface of contact is always 
small during the whole duration of the action. In other 
cases, however, the surface directly acted upon may be large 
as in the case of a vessel driven by sails. In this second case 
it is often convenient to know how much force is applied 
to each unit of area,. 'Phis is found by dividing the whole 
force by the area over which it is distributed, and ihe result- 



4 PHYSICS OF CHEMISTRY. 

ing magnitude is called the pressure. The unit of pressure 
obtains when a force of one dyne is distributed over each 
square centimeter of area. 

Practical Units. — Although the system defined above is the 
simplest and most convenient for all exact calculations, and 
even understanding of complex physical phenomena, it has 
the disadvantage of being in some of its features quite remote 
from our ordinary experience. Hence, although it has made 
its way into the science of physics among all nations, in those 
branches of physical science where only comparatively simple 
conceptions of magnitude, or ratios of magnitudes, are dealt 
with, the need of a logically consistent system of units has not 
been felt. Thus in chemistry, although the rigidly scientific 
system has much to recommend it, its introduction here would 
necessitate so great a change in chemical terminology that the 
loss would be more than the gain. Therefore, although the 
units as defined up to, and including, density are used in 
this text-book, the units of force and pressure, in accordance 
with ordinary usage, are defined otherwise. 

All bodies are attracted towards the earth in the direction 
of the plumb-line. This particular manifestation of force is 
called weight, and for the same body it is found to be nearly 
the same all over the world.* 

Thus the weight of one gram of matter maybe taken as the 
unit of force. The great objection to this definition is that 
the term " gram weight " has two very distinct meanings: it 
may mean a unit by which we measure quantities of matter, or, 
a unit (not perfectly definite) by which we measure quantities 
of force. It is in the former sense that it is generally used, 
and always in this book unless specifically excepted. 

Instead of defining pressure in terms of this new unit of 
force, namely, a "gram weight " per square centimeter, chem- 

* The total variation is about one half of one per cent. 



FORMS OF MATTER. 5 

ists are accustomed to take a unit 13.6 times as large, that is, 
the weight of a cubic centimeter of mercury divided by a 
square centimeter. This practice has grown up from the 
almost universal use of mercury as a means for measuring 
fluid pressures, not only in the barometer for measuring the 
pressure of the atmosphere, but also in other forms of manom- 
eters. This practice has also fixed the habit of designating 
pressures by lengths; thus, a pressure of 76 cm. or 760 mm. 
means a pressure equal to that due to the weight of mercury 
at a surface 76 cm. below the free surface of a vessel of mer- 
cury. To reduce it to grains weight per square centimeter, we 
must multiply the number indicating the length in centi- 
meters by 13.6, the weight in grams of a cubic centimeter of 
mercury. Since nearly all problems in chemistry have to do 
with ratios of pressures only, and the ratios of the lengths of 
the measuring columns of mercury are the same as those of 
the absolute pressures, this process of reduction is rarely 
necessary. 



Forms of Matter. 



Matter appears to us in various forms, which, however, can 
all be reduced to two classes, namely, solids and fluids. 

A Solid is a body which offers resistance both to change of 
shape and to change of bulk. 

A Fluid is a body which offers no resistance to change of 
shape. 

Fluids, again, can be divided into two distinct types: liquids 
and aeriform bodies. 

Liquids are those fluids which resist forces tending to in- 



b PHYSICS OP CHEMISTRY. 

crease their bulk as well as those tending to diminish it. On 
account of this tendency of a liquid to keep its volume un- 
changed^ it is possible to keep a vessel partly full of a liquid, 
in which case the liquid will be bounded by one or more free 
surfaces, i.e., surfaces which separate it from an aeriform 
body. Water is the most familiar example of a liquid. By 
experiment it has been found that a pressure of one atmo- 
sphere (equals 76 X 13.6 grams weight per square centimeter) 
will reduce its volume one twenty-thousandth part only. On 
the other hand, it will resist a force at least ten times as great 
as this which increases its bulk before breaking — a strength 
doubtless greater than that of many friable solids. 

Aeriform Bodies are those which oppose resistance to forces 
tending to diminish their bulk, but none whatever to a force 
tending to increase it. For this reason it is impossible to 
have a vessel partially filled with an aeriform body, the re- 
mainder being empty. Xor, since all aeriform bodies are 
perfectly miscible, is it possible for an aeriform body to have 
a free surface, i.e., to' be self -bounded. Another convenient 
distinction between liquid and aeriform fluids, depending on 
the possibility of a free surface, is this: — a liquid can be poured 
in drops, but an aeriform fluid cannot. 

Gases and Vapors. — Certain aeriform bodies, such as air, 
hydrogen, oxygen, etc., at ordinary temperatures, are subject 
to a very simple law known as Boyle's law, namely, that if the 
temperature remain unchanged the product of the pressure 
and volume is constant. Such bodies are called gases. If 
the aeriform body does not follow this law it is called a vapor. 
We have then the following definitions: 

A gas is an aeriform fluid in which the product of the pres- 
sure by the volume is constant at a constant temperature. 

A vapor is an aeriform fluid in which, at a constant tem- 
perature, this product is not constant. 



THE BALANCE. 7 

It is important to note that experiment proves that every 
vapor becomes a gas at a sufficiently high temperature and 
low pressure, and, conversely, every gas becomes a vapor at 
sufficiently low temperature and high pressure. For ex- 
ample, steam at a pressure of 76 cm. and a temperature of 
200° 0., or at a pressure of 0.1 cm. and a temperature of 20° 
C, may be regarded as a gas; while carbon dioxide is dis- 
tinctly a vapor at a temperature of 30° 0., if the pressure be 
60 or 70 times as great as that of the atmosphere. 

It may be useful to recapitulate the different forms of mat- 
ter with which we have to deal in the science of chemistry. 
They are: 

I. Solids. 

( Liquids. 
II. Fluids. \ . r 

(Aeriform bodies. J <*^ 



The Balance. 



Sir Isaac Newton demonstrated that every body is attracted 
towards the earth by a force dependent directly upon the 
quantity of matter in it, or, in other words, that the weights 
of all bodies are directly proportional to their quantities of 
matter. A balance is an instrument for determining equality 
of weights. It consists of a light rigid beam supported so as 
to turn easily in a vertical plane about its middle point. To 
each end of the beam is suspended a pan, the two being so 
adjusted that when unloaded the beam remains horizontal; 
the position of the beam is determined by a pointer attached 
to it and moving over a graduated plate. If precisely equal 
masses are placed in each pan the tendencyfor each fco descend 
is equal, and the beam remains horizontal; but if one of the 



8 



PHYSICS OF CHEMISTRY. 



masses is greater than the other, that pan in which it is placed, 
on account of its greater weight, will sink. If, then, we have 
a set of standard masses (called iveights) and a balance, we 
may determine the quantity of matter in a body up to the 
value of the sum of the standard masses, and to a degree of 
accuracy limited only by the delicacy of the balance. The 
process of determining masses is called weighing. 




Fig. i. 



A form of balance extensively used is illustrated in Fig. 1. 
It is enclosed in a glazed case so as to protect it from currents 
of air when in use and from dust at all times. Besides the 
essential elements already described, the instrument is pro- 
vided with an arrangement by which the beam can be lifted 
from the delicate knife-edges upon which it is supported, so 
that, when not in use, they are not subject to wear. A por- 
tion of this mechanism can be seen in the figure under the 
ends of the beam: it is actuated by the large knob in front of 
the base of the supporting column. Another important 
accessory is the rod at the top of the case which projects 
through it at the right. By means of this a small bent wire 



THE BALANCE. 



can be placed upon the beam at a distance from its centre 
read from the graduation of the beam. This wire is called a 
rider, and it serves instead of a series of small weights which 
would be inconvenient to handle. 



Determination of Densities. — From the definition of density 
it is only necessary, in order to find it for a given substance, 
to find the mass of a known volume of the substance; then 
the density is equal to the mass as determined by the balance, 
divided by the volume. 

For fluids, the process of determining density is of the 
utmost simplicity. Suppose that we have a thin 
flask, such as is illustrated in Fig. 2, so con- 
structed that when the stopper is in place the 
contents is exactly 100 cc. This condition may 
be brought about by having the glass stopper 
enter the flask so far that the cubic contents is 
slightly too small, and then grinding away the 
lower end" until it is found to contain exactly 
100 grams of cold water. In order to render 
complete filling easy, the stopper has ordinarily 
a fine capillary opening along its axis, and the 
excess of liquid is removed by a bit of filter 
paper so that the surface corresponds to a mark 
on the stem. If now, such a flask be placed 
empty on one pan of the balance and counterpoised, then 
filled with the fluid whose density is desired and again 
balanced, the added weights will measure the mass of 100 
cc. of the fluid, and the density is, if m is the value of the 
added weights, — — - 

° ' 100 cc. 

Example: A flask of 100 cc. capacity was counterpoised on the 
balance; when tilled with alcohol it required added weights equal to 80.95 
grams for equilibrium. From these data we deduce 0.8095 as the 




Fig. 2. 



10 PHYSICS OF CHEMISTEY. 

density of alcohol, i.e., each cc. of alcohol contains 0.8095 of a gram 
of matter. 

The density of a gas is determined in a precisely similar 
manner, although, on account of the relatively small density 
of gases, certain precautions must be taken. For example, a 
flask containing much more than 100 cc. must be employed 
so that its mass may not be too large relatively to that of its 
contents. Again, in this case, when we weigh the flask empty, 
it must be truly void of air as well as all else — a refinement 
which is not essential in tolerably accurate determinations of 
densities of liquids. 

Thus, Eegnault found that a certain glass vessel, by deter- 
mining the mass of water which would completely fill it, con- 
t ined 9881.27 cc. He found also, that when the vessel was 
filled with dry air at a temperature 0° C. and under a pressure 
of 76 cm. it weighed 12.77827 grams more than when quite 
empty; hence the density of air under the conditions given is 
.= 0.00129318 = L nearly. This is the classical deter- 



988] .27 773 

initiation of the important constant, and is the value at present 
accepted as the best attainable. The convenient vulgar frac- 
tion represents the decimal to within less than one part in two 
thousand. 

For solids we must in general take a more indirect way for 
determining their densities. Of course if we have a solid body 
of definita shape and known dimensions — a sphere or a cube, 
for instance — we might readily calculate its volume, and then 
we should only require to determine the mass, and proceed as 
before. But as such a condition of regularity of shape would 
rarely occur in practice, we make use of a principle discovered 
by Archimedes, and known as the 

Principle of Archimedes. — Suppose that we have a vessel of 
water at rest; imagine any sort of closed surface wholly within 







THE BALANCE. 11 

the water, as represented in Fig. 3, separating one portion 

from the rest. It is clear that the water 

outside of the closed surface presses 

upon that within in such a way as to 

keep it in equilibrium, i.e., so that the 

total effect of the pressure from without 

is to sustain the weight of the body of 

water within. Now if the water within 

the imaginary surface be replaced by Fig. 3. 

any other substance, it is manifest that the pressure of the 

water outside would not be modified, and that the resultant 

pressure would still be as before, namely, equal and opposite 

to the weight of the volume of water previously enclosed. 

Hence the principle may be stated : 

If a body be immersed in a fluid at rest, it loses a portion of 
its weight equal to the weight of its own bulk of the fluid. 

The determination of the density of a solid by means of this 
principle becomes easy. The process is as follows : 

Weigh the body in air, then in water; the difference of these 
two weights is the mass of its own bulk of water: divide the 
weight in air by the difference of the two weights and the 
quotient is the ratio of the density of the substance to that of 
water. 

Example : A piece of lead was found to weigh 325 grams in air, and 
296.2 grams in water; hence the loss of weight was 28.8 grams, the ratio 
of its density to that of water -^-g or 11.3. 

Frequently in chemistry we wish to determine the density 
of a substance — a salt, for example — which cannot bo weighed 
in water on account of loss by solution. We may then per- 
form the same operation with some other fluid which does not 
dissolve the substance. The reduction, as above, will yield 
the ratio of the density of the substance to that of the tin id 



12 PHYSICS OF CHEMISTRY. 

chosen, which, multiplied by the density of the fluid, gives 
the density of the material. 

Example: A piece of common salt was found to weigh 172.3 grams 
in air, and 107.2 grams in alcohol, of which the density has been found 
to be 0.8095. Consequently the ratio of the density of the salt to that 



172.3 
65.1 

great, that is, 2.143. 



of the alcohol is ■= = 2.647, and the absolute density 0.8095 times as 

65.1 



Crystallography. 



Only a small portion of the solids brought to the attention 
of the chemist are of that simple character which exhibits like 
properties in all directions with no tendency to assume definite 
geometrical forms. Such simple substances are called amor- 
phous (i.e. , without form). They may be exemplified by glass, 
resins, gums, etc. 

The vast majority of solids, on the other hand, not only 
show marked differences in their physical properties accord- 
ing to the direction along which they are tested, but also a 
remarkable tendency to assume, in the process of forming, 
certain definite geometrical figures bounded by plane faces. 
As an example of physical properties depending upon direc- 
tion we may mention selenite (calcium sulphate), which can be 
broken in three directions only, and with three differing de- 
grees of readiness; or calcite (calcium carbonate), which will 
fracture with equal ease in three directions only, but if heated 
will expand in a certain direction and contract in all directions 
at right angles to it. Again, as an example of the tendency 
to assume definite geometrical forms, we may take common 
salt (sodium chloride), which, if allowed to solidify from an 
aqueous solution, will be found in cubes, or in figures which 
might be built up of cubes. 

A substance which is characterized by such properties is 



CRYSTALLOGRAPHY. 13 

called a crystalline substance. Besides the substances already 
named, we might give ice, most gems, alum, sulphur, sugar, 
etc., etc., as familiar examples of crystalline bodies. 

A piece of crystalline substance formed by continuous and 
spontaneous growth (whether from solution, as salt from brine; 
from fusion, as ice from water; or from vapor, as in the case 
of snow crystals, is immaterial) is called a crystal. The form 
of a crystal of a substance is generally so characteristic that it 
affords a very valuable means of identification. Again, since 
the tendency to crystallize depends largely on the purity of the 
material, crystallization often serves as a means of testing the 
purity as well as a means of separation. Thus an elementary 
knowledge of the classification of crystals is indispensable to 
the chemist. 

In a discussion of the laws governing crystals it is impor- 
tant to observe at the outset that regularity in a crystal is of 
exceptional occurrence. Certain sides may be less favorably 
situated for growth, as, for example, in a crystal lying upon 
one side, that side is obviously unfavorably situated; again, 
in a solution where the temperature is not constant we can 
hardly expect regularity of crystal formation. The elements 
which are constant and characteristic are the angles between 
the bounding planes, and to a less degree the frequency of 
the occurrence of certain faces. 

To make the meaning of the last paragraph clear, we may 
illustrate the first characteristic by the example of crystallized 
quartz. Any crystal of quartz, whatever its color or shape, 
could not fail of recognition if it had a sufficient number and 
variety of faces, for, although a particular angle might be 
found in crystals of other substances, the combination of 
angles proper to quartz does not recur in any other substance, 
The importance of the occurrence of certain faces, the second 
characteristic of a variety of crystals, may be illustrated by 
racemic acid. Of this substance there are two forms which 
have most interesting differences. If we compare the crystals 



14 



PHYSICS OF CHEMISTRY. 



of the two forms we find that each kind possesses faces not 
found on the other, though all the crystal angles are alike in 
the two cases. This peculiarity is the only known method of 
separating the two varieties of the acid. 

The study of the various forms presented by crystals both 
natural and artificial has shown that they may be most simply 
described by reference to certain fixed directions in the sub- 
stance called axes. All forms may thus be referred to one of 
six systems. They are : 

I. Isometric, in which there are three axes at right angles 
to each other, the properties of the substance being alike in 
the three directions. 

II. Tetragonal, in which the properties along two axes 
at right angles to each other are alike, while those along 
the third, which is at right angles to each of the others, are 
different. 

III. Ortliorliomoic, in which again the axes are mutually 




Fig. 4. 



at right angles to each other, but the crystalline properties 
differ iu all three directions. 

IV. Hexagonal, characterized by four axes, three of which 
are in one plane equal and mutually inclined at an angle of 



CRYSTALLOGRAPHY. 



15 



60°, while the fourth, or principal axis, is at right angles to 
the plane of the others. 

V. Monoclinic, in which two of the axes are at right angles 
to each other, and the third, rectangular to one, is oblique to 
the other. 

VI. Triclinia, in which all three axes are oblique to each 
other. 

I. Isometric. — The fundamental principles of crystallogra- 
phy can be illustrated by a consideration of the simpler forms 
of the first class. Thus, in Fig. 4, let oA, oB, and oC be 
the directions of the three axes. Lay off on each equal dis- 
tances from the centre, oa : then the law of crystals states, 
that any plane through the point a which either does not cut 
the other axes at all or cuts them at points a small multiple 
of the distance oa from o, determines the direction of a possi- 
ble pair of faces in the crystal. Examples of forms so derived 
are given in Figs. 5 to 9. 




f^\ 


^ 


i 

1 • 


^* 




i K 





—} 



Fig. 5. Fig. 



Fig. 7. 



Fig. 8. 



Fig. 9. 



The first of the series is the regular octahedron : it is ex- 
emplified in crystals of alum and many diamonds. The last 
is the cube : examples are crystals of common salt, of fiuor 
spar, of iron pyrites, etc. The intermediate £orms. Figs. 6 
to 8, inclusive, possess both octahedral and cubic faces. 

The figures above only represent forms of which the deter- 
mining planes either cut all the axes at equal distances, ov 
cut one axis only. Obviously the system of planes next in 
geometrical simplicity are those which cut two axes only, and 



16 



PHYSICS OF CHEMISTRY. 



those at equal distances. Figs. 10 to 14, inclusive, represent 
derived forms having such planes, together with one or both 
of the other systems, in a greater or less degree of develop- 
ment. The first two of the series represent forms where 
cubic faces and the third class of faces coexist, the cubic faces 
being the more prominent in the first and the latter in the 
second. Fig. 12 represents the form when the planes trun- 
cating the edges of the cube have become so prominent that 
the cubic faces wholly disappear ; it is the geometrical figure 
known as a regular dodecahedron, whence we may properly 
call the faces dodecahedral faces. Crystals of garnet are 
found of this form. Fig. 13 is a form in which all three sys- 




Fig. 10. Fig. 11. Fig. 12. 



Fig. 13. 



Fig. 14. 



terns of faces are represented, namely, the octahedral 1, the 
cubic H, and the dodecahedral L Fig. 11 has only the octa- 
hedral and dodecahedral faces. 

When we consider that this considerable variety of forms is 
derived from determining planes defined by their intercepts 
on the axes, which vary only in the ratio of 1 : 1 or go : 1, and 
that these ratios may be, according to the law of crystal for- 
mation, equal to any number not large,* it becomes obvious 
that the variety of possible forms in this one system is enor- 
mously great. 

One other characteristic of the isometric system should be 
noted, as it distinguishes the first class from each of the oth- 



* A ratio larger than 9 : 1 very rarely occurs. 



CKYSTALLOGKAPHY. 



17 



ers. Inspection of Fig. 5 will show that nine planes can be 
passed through the centre,, dividing the octahedron into sym- 
metrical halves, namely, three defined by each pair of axes, 
and six planes, each of which passes through an axis, and bi- 




Fig. 15. 



sects the angle between the other two axes. These nine 
planes will also divide all the other forms symmetrically; hence 
we say that the isometric system has nine planes of symmetry. 
II. Tetragonal System. — The structure of this system maybe 
elucidated by reference to the crystalline axes, Fig. 15. Here, 




<^ 



Fig. 16. 



^=7* 



Fig. 17. 




^ 



Fig. 18. 



Fig. 19. 



^ 



Fig. 80. 



as before, the axes are mutually rectangular, but we must lay 
off on one axis, called the principal axis, a distance greater 
or less than the distance Oa, laid off on the other two, The 

2 



18 



PHYSICS OF CHEMISTRY. 



ratio of oc to oa, for a given species of crystals must be found 
from a measurement of the angles between the proper faces. 
Having fixed these points, all that has been said concerning 
the derivation of crystalline forms in the first system is applica- 
ble here. Some of the forms are shown in Figs. 16, 17, 18, 




Fig. 21. 

19, 20. Planes perpendicular to the principal axis are called 
basal planes. An example of a crystal in this system is zircon, 
often found of the form of Fig. 20. 

The Tetragonal System has only five planes of symmetry. 
Inspection of Fig. 16 will make this clear. 

III. OrtJ/orJiombic System. — Like the preceding, this is 






Fig. 22. 



Fig. 23. 



Fig. 24. 



characterized by three axes at right angles to each other, all of 
which, however, are of unlike length, as in Fig. 21. The ra- 



CRYSTALLOGRAPHY. 



19 



tios oa : ob : oc in general differ for each species of crystals, 
and can only be determined by measurement of the angles. 
Simple possible forms are represented in Figs. 22, 23, 24, 25, 

26, 27. 




X 



M^=A 



(±=3s kJ^zz^J 



Fig. 25. 



Fig. 



Fig. 27. 



Familiar -examples are Nitre and Aragonite (Fig. 27), Ba- 
rites (Fig. 26), Sulphur (Fig. 25). 




Fig. 28. 



This system has three planes of symmetry only, namely, 
those defined by the axes taken two and two. 



20 



PHYSICS OF CHEMISTKY. 



IV. Hexagonal System. — The characteristics of this system 
are most easily represented, at least geometrically, by reference 
to a system of four axes, three of which lie in one plane and 
are mutually inclined to each other at an angle of 60°, while 
the fourth, called the principal axis, is perpendicular to the 
plane of the others, Fig. 28. The three lateral axes are all 
equal oa, while the ratio of oc to oa is different for different 
species of crystals. Forms derived by the same methods of 
construction as before are shown in Figs. 29, 30, 31, 32, 33. 







i j^i 






\ 






k 


-■ 


- 


.... 


"**v 



f 


? 


\ 


i 




i. 


>. 


/ 




Fig. 29. Fig. 30. 



Fig. 31. 



Fig. 32. Fig. 33. 



The planes perpendicular to the principal axis in Figs. 29, 30, 
32 are called basal planes. In this system the octahedron 
does not occur, but if alternate faces of the double hexagonal 
pyramid Fig. 34 be suppressed, we have a form illustrated in 
Figs. 35, 36 which has, like the cube, only six sides, and is 






Fig. 34. 



Fig. 35. 



Fig. 36. 



called a rhombohedron. Many of the hexagonal crystals — 
calcite for example — will cleave readily in planes parallel to 
these rhombohedral faces. 



CRYSTALLOGRAPHY. 



21 



Examples of crystals of this system are corundum (occurring 
as Fig. 31), quartz (Fig. 33), and calcite (Fig. 36). 

V. Monoclhiic System. — The axial relations are indicated in 
Fig. 37. OA and OB are at right angles to each other, as are 



Fig. 37. 

also OB and 00, but Oa and Oc are inclined. The ratios 
oa : ob : oc may have any value except unity. Simple forms 



are illustrated in Figs. 38 and 39. 






Fig. 



Fig. 39. 



Fig. 40. 



Examples of crystals of this system are feldspar (Fig. 40) : 
and selenite (Fig. 39). 



22 PHYSICS OF CHEMISTRY. 

This system has only one plane of symmetry, namely, that 
of the axes OA and OC. 

VI. Tridinic System. — This is characterized hy three axes, 
no two of which are at right angles, and the axial ratios are 
all unequal. A crystal in this system has no plane of sym- 
metry. The most familiar crystal of this system is that of 
copper sulphate. 

Iso-, Di-, and Pleo-morphism. — Many chemical compounds of 
dissimilar but analogous composition present identical crystal 
forms. Such bodies are said to be isomorphous. As exam- 
ples we may cite the various alums, which belong to the First 
System, and the calcite grou]3, including calcite, dolomite, 
magnesite, etc., which belong to the Fourth System. 

On the other hand, some substances are found to occur in 
crystals belonging to two or more of the six crystalline sys- 
tems. If in two systems only, the substance is said to be 
dimorphous; if in more than two, pleomorphous. A good 
example of a dimorphous substance is sulphur, which, when 
crystallized from the liquid condition, forms crystals belong- 
ing to the Fifth System, but when deposited from solutions 
or from the vaporous state forms crystals of the Third Sys- 
tem. Calcium carbonate is another common example. The 
substance is known under the names calcite, in which form it 
belongs to the Fourth System ; and aragonite, the crystals 
being orthorhombic. 

Molecular Structure of Matter. — In the study of crystals, by 
far the most important fact discovered is that the ratios for 
the interceps on the axes for a possible crystalline face are 
always rational, as defined in the law on p. 15. To interpret 
its physical meaning, let OA and OB (Fig. 41) be two axes of 
a crystal, and Oa x , Oh } represent the interceps of that plane 
parallel to the third axis, which is found to bear the simplest 
relations to all other observed planes parallel to the third 
axis; then the law says that any plane passing through a x and 



CKYSTALLOGRAPHY. 23 

b i9 a x and b 3 , 0, and b i9 etc., or a 2 and 5„ a 3 and b lf and a 4 and 
b xi etc., are possible faces of a crystal, but no others can be. 
But if the material of the crystal were absolutely uniformly 
distributed, there is no assignable reason why a plane other 
than the ones thus defined should not occur. Hence we are 
obliged to conclude that the matter is arranged in isolated 
particles similarly to the arrangement of the dots in the 
figure. With this hypothesis as to the arrangement of the 
matter, it is easy to see what the crystallographic law means. 
The plane a y b x contains every particle of matter in its own 
direction, the planes ajb t and ajb l every second particle, etc., 
etc. This structure, so definitely indicated by crystals, is 




Fig 41. 

called a molecular structure, and the particles are called 
molecules. Beyond this, crystallography teaches us nothing, 
except that structure must be exceedingly fine-grained, for in 
some cases perfectly polished planes containing every fifteenth 
row of molecules have been observed; but if this distance 
were any considerable part of a wave-length of light, i.e.. 
about rfojf mm., the polish could not be perfect: whence 
we conclude from such phenomena alone that there cannot 
be less than fifteen hundred thousand molecules side by side 
in the length of an inch. We shall find more accurate meth- 
ods of indicating their distances and sizes further on. 



24 PHYSICS OF CHEMISTRY. 



Temperature and Heat. 

Difference of Temperature. — When two bodies are brought 
into contact we find in general that one of them becomes 
cooler and the other warmer. The body which becomes 
cooler is said to be at a higher temperature than the other, 
and, other things being equal, we define the difference of tem- 
perature as proportional to the rate at which the process goes 
on. Equality of temperature obtains when neither of the 
bodies becomes warmer or cooler. 

To measure differences of temperature we must take two 
substances having constant differences of temperature and 
give numerical values to each. For example, it is found that 
the difference in temperature between melting ice (or freezing 
water) and water boiling under a fixed pressure is constant ; 
hence, if we call the temperature of melting ice zero degrees, 
and that of boiling water, or more accurately, steam from 
water boiling under the normal atmospheric pressure, one 
hundred degrees, we shall have a scale for measuring differ- 
ences of temperature. Thus, a body which is found to be in 
thermal equilibrium (i.e., neither grows cooler nor warmer 
when immersed in it) with a mixture of equal parts of boiling 
water and ice-water has a temperature fifty degrees higher 
than that of ice, or, in short, a temperature of 50°. Again, 
a body in thermal equilibrium, with a mixture of one fourth 
boiling water and three fourths ice-water, would have a tem- 
perature of 25°. The temperature of a body colder than ice 
would obviously have to be indicated by a negative number. 
It is important to observe that the numerical value of a tem- 
perature is merely the value of the difference of temperature 
between the substance and that of melting ice, measured in 
such units that 100 of them express the difference in temper- 
ature between ice and boiling water. Hence a tenrperature 
of 50° is in no sense twice as high as a temperature of 25°, 



TEMPERATURE AND HEAT. 25 

nor, from what appears in the definition, is there any numer- 
ical ratio between them. 

Thermometers. — The method of determining differences of 
temperature indicated above is not a convenient one ; so, in 
practice, the variation in volume of some substance which is 
found to follow sufficiently closely the same rate as the varia- 
tions of temperature is chosen. Of the various possible 
substances the only one which is generally used in 
chemical work is mercury, which, enclosed in a bulb 
provided with a tube of fine bore, constitutes the 
mercurial thermometer. Fig. 42 shows the form or- 
dinarily used in the laboratory. Three processes are 
necessary to fit it for use. First, when sufficient mer- 
cury has been introduced into the bulb, the top of 
the tube is sealed after excluding the air by boiling 
the mercury; then the bulb and stem up to the top 
of the column are immersed in moist pounded ice 
until there is no further change and the position of 
the top of the column marked on the glass: this is the 
zero point. Second, the bulb and stem are immersed 
in steam issuing from boiling water, when the pres- 
sure of the atmosphere is 760 mm., and the position of 
the end of the column marked on the tube, giving 
the 100° point. Third, the stem is divided between 
these fixed points into one hundred parts of equal 
volume, and, if required, the division is carried on 
above and below the two points. In practice it is 
assumed that equal distances on the stem correspond 
to equal volumes, an approximation which is gener- 
ally sufficient for carefully chosen tubes. 

The thermometer so divided is universally used 
in chemistry, and is known as the centigrade ther- 
mometer. When there is any danger of eon fusing Fu; <. : 
it with any other thermoinetric scale, it is the custom 



25 PHYSICS OF CHEMISTRY. 

to write a capital C after the number expressing the read- 
ing; thus 30 c C. is to be read, thirty degrees centigrade. 

Calibration of Thermometer. — If, however, a more accurate knowledge 
of the thermometer is required, it will be necessary to calibrate the tube 
so as to determine the errors of such a number of points intermediate be- 
tween 0° and 100°, that the errors for every portion may be safely in- 
ferred. The process is as follows. Having inverted the thermometer, 
tap it lightly against the lower end. Then either the thread will sepa- 
rate, or the mercury in the bulb will fill the tube, leaving a small space 
at the base of the bulb. If the latter is the case, the thermometer must 
be turned and the vacant space brought to occupy the region of the 
opening of the bulb — a condition which is always practicable. We thus 
secure finally a thread of mercury in the stem separated from the re- 
mainder by a vacant space. The length of this thread can be modified 
at will. Suppose, for example, we wish to make it n degrees longer. Erect 
the thermometer and allow the thread to run down to contact with 
the rest; then warm the thermometer n degrees, invert and tap again, and 
the thread will break at the same point as before, leaving thus an 
increased length of the desired amount. If, on the other hand, we 
wish a thread shorter by n degrees, we cool the bulb after the thread has 
been allowed to run down into contact with the remainder by this 
amount, and produce a separation again as before. The fracture of the 
thread is determined by the presence of a minute bubble of air which 
clings to the wall of the tube . and does not move with the mer- 
cury. It is obvious that in this process the bulb should be warmed suf- 
ficiently to have the point of contact in the tube and not in the bulb. 

Having thus secured a thread as nearly as may be of 50° in length, 
cause it to slide along the tube until one end is at 0° and the other, say, 
at 50°. 6. Then shift the thread until the upper end is at 100°, the lower 
reading, say. 49°. 7. Since the volume ( F) of the mercury in the thread 
remains unchanged, this observation gives us 

50 

Volume from to 50 = —-- V, 
o0.6 

50 
Volume from 50 to 100 = =jt-= V, 

oO. o 

50 3 
whence the ratio of the first volume to the second is — -r, and, since the 

o0.6 

temperature interval from the 0° point of the scale to the 100° point has 



TEMPERATURE AND HEAT. 27 

been determined by ice and boiling water, we have the temperature in- 
terval between the 0° point and the 50° point 

100.9 ' 

if T is the former interval. Thus the true temperature corresponding 
to the thermometric reading 50° is :~ T-\- k , if k is the correction to 
be added to the zero reading of the thermometer. 

With a thread of nearly 25° in length it will be easy to determine 
by a precisely similar process the errors of the 25° and 75° points after 
having determined that of the 50° point, and so on for any number of 
aliquot subdivisions of the scale desired. In practice it will be found 
sufficient to determine these five points unless the errors are found 
to be considerable, in which case the thermometer should be rejected 
for all exact use. 

Heat. — When equal quantities of two unlike substances at 
different temperatures are kept in contact until these temper- 
atures are alike, it is found that the final temperature is in 
general very, far from the average of the two. Since in this 
case we suppose that all the heat which has left the hotter 
body goes into the colder, we. must conclude that it requires 
very different amounts of heat to raise like masses of dif- 
ferent substances through the same range of temperature. 

Exp. 1. — This important fact can be demonstrated qualitatively by 
placing a beaker of water over a Bunsen burner and noting the rate 
of the rise of temperature by means of a thermometer; then replace 
the water by an equal amount of some other liquid, say turpentine, 
and again note the rate of increase of temperature: it will be found 
to be much greater in the second case, more than twice as great if tur- 
pentine be used. As the source of heat is the same in the two cases, 
and all the other conditions are very nearly the same, we conclude that it 
requires more than twice the quantity of heat to raise a given quantity 
of water through a given range of temperature, than it does to raise a 
like quantity of turpentine through the same range. 

Unit of Heat: Calorie. — In order to define a quantity of 



28 PHYSICS OF CHEMISTRY. 

heat we must therefore not only consider the change of tem- 
perature which it produces, but the magnitude and substance 
of the body in which the change is produced. The heat unit 
employed in chemistry is a quantity of heat which added to a 
gram of water at 0° C. would raise its temperature 1° 0. This 
unit is called the calorie. 

Specific Heat. — The specific heat of a substance is the ratio 
of the quantity of heat required to raise the temperature of 
one gram by 1° C. to the calorie. 

Determination of Specific Heat. — We shall consider here 
only the method known as the method of mixture. Suppose 
we have a vessel (of thin metal by preference) containing M 
grams of water at a temperature t, into which we put m grams 
of the substance whose specific heat we wish to determine, at 
a higher temperature f. The substance may be liquid or 
solid, but if the latter it should be in small pieces. After the 
temperature of the water and the substance has become the 
same, which can be hastened by stirring, read the temperature 
with a thermometer; call it t" . Then the rise of temperature 
of the water is t n ' — t, and the fall of temperature in the sub- 
stance is f — t"\ consequently the water has gained M{t" — t) 
calories, and, if s be the specific heat of the substance, the 
amount of heat lost by it is equal to ms(f — t") calories. 
These two quantities must be equal if precautions have been 
taken against gain or loss from outside, hence 

_ M (t " - t) 
S ~ m (? - t") 

Fusion. — "When heat is added continuously to a solid its 
temperature rises continuously up to a certain point, when it 
commences to melt. If the mixture of the solid and liquid is 
constantly stirred, it is found that the temperature remains 
unchanged until it becomes entirely liquid. This fixed tern- 



TEMPERATURE AND HEAT. 29 

perature is called the melting point, and it is quite character- 
istic of a substance. The following table gives a number of 
melting-points. The higher temperatures are only approxi- 
mate. 

Table of Melting Points. 

Mercury -39° Gold 1240° 

Tin 228° Copper 1830° 

Lead 334° Cast-iron 1200° 

Zinc 423° Wrought-iron ' 1600° 

Silver 1000° Platinum 2000° 

Vaporization. — If heat is added continuously to the bottom 
of a liquid, the temperature rises until bubbles of vapor form 
at the lower surface and rise through the body of the liquid. 
This phenomenon is called boiling, and it is found that the 
temperature remains unchanged until all the liquid has been 
converted into vapor. The temperature of boiling is always 
the same for the same liquid, provided that the pressure is 
the same; but it varies greatly with the pressure. By the 
boiling point is meant the temperature of a boiling liquid 
under the standard atmospheric pressure, namely, 760 mm., 
mercury column. The following table contains some of the 
more important boiling points. The very low and very high 
boiling points are only approximate. 

Table of Boiling Points. 

Hydrogen -215° Water 100° 

Oxygen — 181° Mercury 350° 

Carbon dioxide — 78° Sulphur 448" 

Ammonia — 34° Zinc 1040° 

Sulphur dioxide — 1(P Silver Oxyhydrogen dame. 

Ether 35° Platinum and other difficult- 
Alcohol 78° ly volatile metals.. In voltaic arc. 



Latent Heat of Fusion. — To convert a given quantity of a 

solid into a liquid at the same temperature requires the addi- 



30 



PHYSICS OB CHEMISTRY. 



tion of a large quantity of heat. Since the addition of the 
heat is unaccompanied by rise of temperature the heat thus 
expended is called latent heat. 

Exp. 2.— Mix 150 grams of water al 100' with 100 grams of ice at 
: . The ice will be melted and the temperature of the mixture will 
be about 28\3 C. exactly if the experiment be very accurate;. In 
this case the water has lost 10,755 calories., and the water of the melted 
ice has gained "2880 calories. The difference is 7925 calories, which has 
been employed in converting the 100 grams of ice at : into water at r ; 
hence it requires 79.25 calories to convert a gram of ice into a gram of 
water at the same temperature. This number is called the latent heat 
of fusion of ice. 






Bunsen's Ice Calorimeter — If a body at a temperature, say. 
of 100 z C. is placed in a cavity in a block of ice at 0° C. it will 
finally attain the same temperature as the ice. after having 
melted a certain number of grams. The number of grams 
of water from the melting multiplied by 79.25 will give 
the number of calories lost by the body, 
C which, divided by the product of the 

mass of the body by the product of its 
fall in temperature, gives the numerical 
value of the specific heat. This principle 
is the basis of Bunsen's calorimeter, an 
apparatus which has proved of great 
value to chemistry, because it admits of 
accurate determination of the specific 
heat of a very small quantity of a sub- 
stance. It consists of a glass vessel. A 
(Fig. 43), having sealed into it a closed 
tube B. and a fine graduated tube C. 
The tube C. together with a portion of 
the vessel A, is partially filled with mer- 
FlG 43 cury. the remaining portion of A which 

surrounds B beins: filled with water carefully freed from air. 




PRESSURE AND VOLUME OF OASES. 31 

Then a current of alcohol at a temperature below zero is 
passed through the tube B until all of the water in A is 
frozen. Since ice is nearly one tenth more bulky than the 
water from which it is derived, a portion of the mercury will 
escape from C during this process. Next, the whole apparatus 
being buried in moist snow, the temperature throughout 
finally becomes zero, and the end of the thread of mercury 
assumes a fixed position in C. Let this position in the gradu- 
ated tube be designated by a. Now, after having removed 
the cold alcohol, drop a gram of water at 100° into the tube 
B, a certain portion of the ice will be melted, and the end of 
the thread of mercury will be retracted, reaching finally, after 
the water has cooled to 0°, a position on the scale b. Thus a 
change of b — a divisions on the scale corresponds to an addi- 
tion of 100 calories to the calorimeter. If now a gram of the 
substance to be tested at a temperature of 100° is dropped into 
the tube B, more ice will be melted and the thread will be 
farther retracted to the point c; consequently the ratio of the 
heat lost by -the substance to that lost by the gram of water is 
c — b : b — a; but this is obviously the specific heat of the 
substance, as appears from the definition. Since the tempera- 
ture of 100° can be accurately secured by keeping the substance 
in a vessel immersed in steam from boiling water sufficiently long 
it is obvious that no thermometer is required in the experiment. 



Pressure and Yolume of Gases. 

Barometer. — The pressure of the atmosphere can always be 
determined by the barometer. This, in the form used in the 
laboratory, consists of a glass tube closed at the upper end and 
bent in U form near the lower end, as in Fig. 44. The distance 
from the bend to the closed end must be considerably greater 



32 



PHYSICS OF CHEMISTRY. 




than 76 cm. Into the tube, when held in a reversed position 
from that in the figure, are introduced suc- 
cessive small quantities of mercury, which are 
boiled over a lamp so as to expel all the air 
clinging to the glass and to the surface of the 
mercury. When this process has been car- 
ried on until the tube is full to the bend, it 
is erected and the mercury falls a certain 
distance in the long arm, leaving a vacuum 
above. The difference of level in the free 
surfaces of the mercury is determined from 
millimeter graduations on both arms of the 
tube, and is found 
to be at the sea- 
level about 760 mm. 
Since the atmos- 
pheric pressure is 
cut off from the 
upper end of the 
fig. 44. tube, it is obvious 

that the pressure of the air on the 
free surface of the mercury in the 
short tube is equal to the pressure 
due to the weight of a layer of mer- 
cury whose depth is equal to the 
difference of level in the two arms. 
The difference 760 mm. is taken as 
the measure of the standard at- 
mospheric pressure. This arbitrary 
unit of pressure is designated by the 
term at mo. 

Fig. 45. 

Measurement of Gases. — By means 
of a barometer and a graduated tube closed at one end, it 
is easy to measure the volume and pressure of a quantity of 




LAWS OF GASES. 33 

gas. Suppose the tube to be graduated to cubic centimeters 
from the closed end. If it be filled with mercury, inverted 
with its open end under the surface of a vessel of mercury, 
and a sufficient quantity of gas be then introduced into the 
tube, it will assume the appearance of Fig. 45. The volume 
of the gas can be read at once from the tube, while its pressure 
is equal to that of the external atmosphere diminished by 
that due to a column of mercury of height x. The pressure 
is then, if li is the observed height of the barometer, 

— — — X atmo. , where all the lengths are measured in milli- 

meters. A thermometer suspended in contact with the tube 
will give the temperature of the gas. 



Laws of Gases. 



The law connecting pressure and volume of a gas at con- 
stant temperature, known as Boyle's law, has already been 
given. It may be formulated thus : 

pv t = h [t constant], 

where p equals pressure, v t the volume at temperature /, 
and k a constant. 

The law connecting volume and temperature at constant 
pressure is equally simple. It has been found that all gases 
expand ¥ -£ ¥ of their volume at 0° C, for each increase of one 
degree in temperature, provided that the pressure remains 
unchanged. This law, discovered by Charles but commonly 
known as that of Gay-Lussac, may be formulated thus : 

Vt = v (l + *Jy) [p constant], 

where v t equals the volume of the gas at the temperature /, 
v the volume at 0° C, / the temperature on the centigrade 
scale, and p the pressure. 
3 



34 PHYSICS OF CHEMISTRY. 

These two equations can be combined so as to yield a single 
equation applicable to all cases : for let P represent the 
standard atmospheric pressure and V the volume of the gas 
under this pressure and at a temperature of zero, then the 
first equation gives 

pv =PV 9 = l\ 

and the second gives a value for v , which, substituted here, 
yields 

PV o=P 1 + i o 036T ^ 

where the value of the fraction ^3- is expressed by the 
decimal .00367. Thus, if we observe the volume v t of a gas 
under a pressure p and at a temperature t, we can find from 
this formula what its volume should be at a temperature 0° 
and at atmospheric pressure even if the substance in question 
cannot exist as a gaseous body at that pressure and tempera- 
ture. We may make a convenient modification of this for- 
mula to adapt it to practice. Since gaseous pressures are in 
direct ratio to the heights of the columns of mercury which 
they will sustain, and the normal atmospheric pressure is 
assumed to be equal to that of a column of mercury 760 mm. 
in height, we may write 

ll Vt 



760 mm 1 + .00367*' 



where li is to be measured in millimeter: 



Kinetic Theory of Gases. 

Kinetic Theory. — A gas has been defined as an aeriform 
body which obeys Boyle's Law. i.e.. as one in which the pres- 
sure at constant temperature is inversely as the volume. 
Hence if the pressure be diminished without limit the volume 
would increase without limit, or. in other words, the mole- 



KINETIC THEORY OF GASES. 35 

cules would separate indefinitely from each other. The most 
obvious conclusion is that the molecules repel each other; but 
Daniel Bernoulli pointed out before the middle of the last 
century that it is impossible to assign any law of repulsion 
which would make the pressure independent of the shape and 
size of the containing vessel — in short, which would make the 
pressure follow the well-known law. He showed, also, that if 
all the particles be regarded as in motion, the pressure might 
be explained by the collision of the molecules against the sides 
of the vessel. This general explanation has received a wide 
development during the last forty years, under the name of 
the Kinetic Theory of Gases. 

Explanation of Gaseous Pressure. — The fundamental sup- 
positions of the kinetic theory are, first, that every body of 
gas which can be experimented upon contains a practically 
infinite number of molecules ; second, that all the mole- 
cules are in continuous motion with a high velocity ; third, 
that the dimensions of the molecules are very small compared 
to their average distance apart ; and, fourth, that the time 
during which any one molecule is near enough to any other 
to act upon it is very brief compared to the time that it is too 
remote from all others to affect them. 

With these suppositions it is easy to find an expression for 
the pressure of a gas in terms of its density and the average 
velocity of its molecules. 

Let Fig. 46 represent a rectangular box having the edges 
a, b, c, filled with gas. Let 
p = pressure of the gas. 
N =■ number of molecules in one cc, 
m — mass of each molecule, 
p — Ntn = density of gas, 
V= average velocity of motion of mole- fig. 46. 

cules. 

If a molecule strikes normally against a side of the vessel; 



& i 

c.- ■•■' 

a 



36 PHYSICS OF CHEMISTRY. 

say the side be, it will, on the average, bound off with the 

same velocity with which it struck ; the motion is thus 

changed from mV to — m V, a total change of 2m V, by the 

reaction of the wall. Since force is measured by the rate at 

which it changes motion, the average force exerted by the 

2m V 
wall upon the molecule is , if r is the duration of the 

whole action. Now by the second supposition the number of 
molecules, and hence the number of molecular impacts per 
second, is indefinitely great; therefore the force exerted by the 

2m V 

wall is constant, and equal to multiplied by the number 

of impacts in the time r. To find the whole constant force 
upon be, then, we have only to find the number of molecules 
which strike against it during the time r. 

It is obvious that we shall not be very far from the truth in 
calculating the number of impacts if we make the simple sup- 
positions, first, that all the molecules have the constant ve- 
locity of the average, and, second, that a third of the mole- 
cules are moving parallel to each of the three edges a, b, and e. 

The first supposition gives the number of times which any 
one molecule of those which move parallel to a strikes the 

V V 

side be in a second as — -. or during the time r as — r - This, 
2a 5 2a 

with the second supposition, since the whole number of mole- 
cules in the box is Ndbc, gives 

1 VI 

—]\ abc-—T = —Nbc Vr 



as the number of molecular impacts on be during the time r. 
We thus have the total constant force exerted by the face be, 

F=\KmbcV\ 

o 



KINETIC THEORY OF GASES. 37 

But pressure is defined as force divided by area upon which it 
is exerted; hence 

But this is obviously the law of Boyle. 

Calculation of Molecular Velocities. — From the last equa- 
tion we have 



T= \l\ 



In order to reduce this we must express p in the same units 
as are chosen for V and p, namely, in centimeters, grams, and 
seconds. As a special example, we will find the molecular 
velocity for air at 0°C. and 76 cm. pressure, measured in mer- 
cury column. 

The pressure at the base of a mercury column 76 cm. in 
height is 76 X 13.6 grams weight per square centimeter. But 
if a gram be allowed to fall freely it is found that during the 
first second it will have acquired a velocity of 980 cm. per 
second, that is, a gram weight is 980 times the unit of force 
defined in terms of the gram, centimeter, and second, or 980 
dynes. Again, experiment, such as described on p. 10, shows 
that the density of air at 0° and atmospheric pressure is ^-3, 
hence, substituting these values, 



V = Vd • 76 • 13.6 • 980 • 773 = 48500 cm. 

Again, since the molecular velocity for a gas at 0° varies 
inversely as the square root of the density, we may find the 
velocity for any other gas by multiplying this number by the 
square root of the ratio of the density of air to that of the gas 
in question ; e.g., air is 14.44 times as dense as hydrogen. 
hence the mean value for the velocity of a hydrogen molecule 
at a temperature of 0° is 48500 cm. X ^U.44 = L84300 centi- 
meters per second, approximately 1 mile per second. 



L698 meters. 


Carbon dioxide, . 


. 361 meters. 


566 " 


Chlorine, . . . 


. 286 " 


453 " 


Bromine, . . . 


.190 " 


447 " 


Mercury, . . . 


. 169 " 


425 " 


Iodine, .... 


. 151 " 



38 PHYSICS OF CHEMISTRY. 

In the above calculation the supposition which is most 
questionable, and which would give rise to the greatest errors 
in the results, is the fifst, namely, that the molecular veloci- 
ties are all equal. For even if this equality of velocities ex- 
isted for a moment it would be immediately destroyed by 
mutual molecular collisions, and inequality follow. Taking 
into account, however, this fact of constantly varying veloci- 
ties, Maxwell has found that the average velocities are || of 
these, calculated according to the simple methods above. 
Thus corrected, the mean molecular velocities for a number of 
substances at 0° is given below: 

Hydrogen, . 
Water vapor, 
Nitrogen, . 
Air, . . . 
Oxygen, . 

Law of Avogadro. — The most important result of Maxwell's 
investigations into the laws of molecular motion in a gas was 
the proof that in any two gases at the same temperature the 
products of the molecular mass into the square of the average 
molecular velocity are equal. Hence, if we have two gases, 1 
and 2 at the same temperature, 

Also, if the pressures be the same, N l m l V* = iV 2 m 2 F 2 2 , each of 
these two equal quantities being three times the pressure. 
Dividing the second quantities by the first, we have 

x, = x. ; 

that is, In equal volumes of any two gases at the same pres- 
sure and temperature there are an equal number of molecules. 
This law, of the highest value in chemistry, is known as the 
Law of Avogadro ; it enables us to determine the relative 
masses of the molecules of all substances which can be experi- 
mented upon in the gaseous condition. 



KINETIC THEORY OF GASES. 



39 



Diffusion. — Suppose the box (Fig. 46, p. 35) to have a small 
opening of area s in the end be ; then, as appears from the 
reasoning on p. 36, the number of Inolecules which would 
reach the aperture per second and escape from the box would 
be -j-iVs V. For all gases, however, pressure and tempera- 
ture being the same, N is the same ; hence the rate at which 
the gas would escape would be proportional to V, or inversely 
proportional to the square root of the density, since V for dif- 
ferent gases varies in this ratio. This conclusion is found to 
be in accordance with experience. 

Again, suppose the box divided into two sections by a por- 
ous partition, say of plaster of Paris or of unglazed earthen- 
ware, which may be regarded as a wall with innumerable 
small openings in it. If the sec- 
tions be filled with gases at like 
pressures and temperatures, but 
of unlike densities, and conse- 
quently of unlike molecular veloci- 
ties, more x>i the lighter molecules 
will penetrate the porous partition 
in a given time than of the denser, 
hence the pressure will at first in- 
crease on the side of the denser 
gas, and decrease on the other. 
Finally, however, the mixtures 
will become alike on the two sides 
of the diaphragm. 




Fig. 47. 



Exp. 3. Diffusion of hydrogen. — Fig. 
47 represents two bell- jars, the upper 
one of which is closed at its base by a 
plate of plaster of Paris, each having a manometer tube attached. The 
bases of the jars are ground true, and greased so as to tit together gas 
tight. 

To make the experiment, remove the manometer tubes and place the 
upper jar on a glass plate, and rill it with hydrogen by passing a rapid 
current of the gas into the jar through the lower tubulure. Next. 



40 PHYSICS OF CHEMISTRY. 

place the jars in position as shown in the figure, and quickly adjust 
both manometer tubes at the same instant. 

The manometers will show a marked decrease of pressure in the 
upper jar and an equal increase in the lower. After some time this 
difference of pressure will attain a maximum, and then slowly disappear. 

AYe can draw a conclusion of great interest to the science of 
chemistry from an experiment of this kind. Suppose in the 
box, arranged as last described, one of its divisions only be 
filled with a gas, e.g., hydrogen, the other division being left 
empty ; then, if hydrogen consists of molecules of different 
masses, since the lighter ones must have a higher average ve- 
locity in accordance with Maxwell's law, we should be able to 
sift out the lighter molecules and thus secure a hydrogen of less 
density than that with which we started. As such separation 
is found to be impossible, we conclude that all the molecules 
of a gas are alike. 

Size and Number of Molecules. — The average distance that 
a gaseous molecule will move before coming into a collision 
with another is called the mean free path. This depends evi- 
dently upon the average space occupied by each molecule, and 
the cross-section of the molecule ; for if either the density of 
the gas or cross-section be diminished the relative portion of 
the space occupied by matter will be reduced, and the chance 
of a molecule going a longer distance before meeting another 
will be increased. The law connecting these three quantities 
has been deduced from the kinetic theory of gases, so that if 
any two of them are known from observation the third can be 
calculated. But the mean free path can be derived from ob- 
servations on the viscosity of gases with considerable precision. 
On the other hand, although we have no means of finding ex- 
actly the cross-section of the molecules, still the great incom- 
pressibility of liquids suggests that the molecules are probably 
not far removed from contact : hence if the volume of a gas 
when condensed to the liquid form is known, we may estimate 
the sum of the cross-sections of all the molecules contained in 



DETEBMIKATION OF GAS DENSITIES. 41 

a unit volume of the gas. These two data yield a value for 
the number of molecules in unit volume of the gas, and there- 
fore the approximate size also. This process applied to the 
gases carbon dioxide, hydrogen, and air give diameters 0.18, 
0.14, and 0.30 millionths of a millimeter, respectively. These 
values, all of the same order of magnitude, do not differ great- 
ly from other estimates made by entirely independent meth- 
ods, and are thus entitled to considerable confidence. It is 
certainly hardly possible that they are either ten times too 
great or ten times too small. 

The same process yields the number of molecules in a cubic 
centimeter of air, or, in accordance with Avogadro's law, in a 
cubic centimeter of every gas. It is, for atmospheric pres- 
sure and 0° C, 

21(10 18 ), 

or, 21 millions of millions of millions. Since we cannot ob- 
serve a gas under a pressure many hundreds of times less than 
the atmospheric pressure, it is obvious that we are always jus- 
tified in assuming that every body of gas observed contains 
practically an infinite number of molecules. 



Determination of Gas Densities. 

The important principle stated on p. 38, and known as 
Avogadro's law, enables us to determine the ratio of the mass 
of any gaseous molecule to that of a molecule of hydrogen if 
we know the ratio of the density of the gas to that of hydrogen 
at the same temperature and pressure. The determination of 
this density ratio, commonly called the gas density or vapor 
density, becomes thus a process of fundamental importance in 
chemistry. There are many ways* of doing this, each of which 
possesses advantages of its own; but the method devised by 
Victor Meyer for determining the gas densities of substances 
which are ordinarily in the solid or Liquid state, but which may 



42 



PHYSICS OF CHEMISTRY. 



be made to assume the gaseous condition at a temperature 
below the point at which glass softens, and at a moderate pres- 
sure, is of extraordinary simplicity. 

The apparatus employed is shown in Fig. 48, a. It consists 
of a glass cylindrical vessel B, of about 
100 cc. contents and a length of 200 
mm., to which is sealed a tube of 
about 600 mm. in length and 6 mm. 
in diameter, ending above in an' 
enlarged mouth stopped by a rubber 
stopper. A little below the upper 
end is a small branch tube A, the 
end of which, bent upwards, is below 
the surface of a vessel of water D. 
If a temperature not exceeding 310°, 
the boiling-point of diphenylamin, is 
sufficient, the tube B is surrounded 
by a larger glass tube ending in a bulb 
C, which contains the liquid for heat- 
ing the tube B. This liquid may be 
water (to 100°), aniline (to 183°), 
amylbenzoate (to 260°), or diphenyl- 
amin (to 310°). For still higher 
temperatures the vessel B is intro- 
duced into a piece of closed iron gas- 
pipe M, which is surrounded by melted 
lead contained in the larger pipe S, 
Fig. 48, I. 

To determine the gas density of a 
substance a quantity is weighed out, 
which, when gasified, will not more 
than half fill the vessel B. This is 
enclosed in a little glass flask suffi- 
ciently small to drop through the 
tube into B. After the apparatus has attained a constant 




DETERMINATION OF GAS DENSITIES. 43 

temperature so that there is no further escape of air through 
the tube A, the stopper is removed, the substance dropped 
into B, where there is a cushion of asbestus to protect the 
glass, and the stopper immediately replaced. As this is 
pushed in one or two bubbles of air will escape from A, after 
which the inverted graduated tube, quite full of water, is to 
be placed over the end of A . The substance will evaporate 
in a few seconds, and its gas will drive out its own bulk of air 
which will collect in the graduated tube. After the flow of 
air has ceased, the graduated tube is lowered in the water until 
the level of the surface of the water inside and outside is the 
same, and the volume of the air is read from the graduation, 
together with its temperature, by means of a thermometer. 
This, with a reading of the height of the barometer, completes 
the observations. 

The reduction of the observation is as follows. It is obvious 
that the mass of the gas formed from the substance is the 
same as the mass of the substance, and is consequently knowm. 
Let this mass be m. The mass of the air displaced may be 
represented by m' , and its observed volume by v. Since the 
density of air is 14.44 times as great as that of hydrogen, w r e 
have the gas density of the substance, as defined above, 

771 

equal to D = — 7 14.44. 

^ 71V 

It only remains to find the value of m'. As appears on p. 
10, the mass of a cubic centimeter of air at a temperature 0° 
and a barometric pressure of 760 mm. is T 4~3 of a gram ; but 
by the law of Boyle and Charles, p. 34, we can find the vol- 
ume reduced to this standard temperature and pressure, when 
the volume, temperature, and pressure are known. Thus, 

T . h v± 



760 mm. (1 -f .00367/)' 
and 

m' = ^ grams, 
776 



44 PHYSICS OF CHEMISTKY. 

In these formulas V is the reduced volume, v t the observed 
volume, h the barometric reading at the time of the experi- 
ment, and t the temperature. 

This process would be complete if the displaced air were 
collected over mercury, but as it is more convenient to collect 
over water, a connection is required for the vapor of water, 
which is mixed with the air in the graduated tube. As the 
air has not only passed through the water in small bubbles, 
but has stood in contact with it for some time, we may safely 
assume that at the time of reading the volume the pressure 
of the water vapor is as great as possible under the existing 
condition of temperature, and can be found from the table 
of maximum pressures of water vapor given on page 113, cor- 
responding to the observed temperature. Let this pressure be 
denoted by w. It is obvious that the whole pressure in the 
graduated tube, measured by h, is the sum of the pressures of 
the air and of the water vapor; hence the pressure of the air 
is measured by A — w, which is to be substituted in the equation 
above. Making this substitution, we have 

h — w v t 



760 1 + .00367T 
y 

m t _ _l_o_. 
/ (6 

and finally, 

m _ 14.44m • 760 • 773(1 + .00367*) 
m (h — w)v t 

Carrying out the multiplications of the constants, we may 
write 

m(l + .00367*) 



D = 8484000 



(Ii — w) v 



The results of some of Meyer's determinations, and a com- 
parison with the values calculated from theory, according to 
methods given farther on, may be given here. 



DETERMINATION OF GAS DENSITIES. 45 

Chloroform [heated in steam]. 

m = 0.1008 t = l6°.5 li = 707.5 v = 22 cc. 

D 59.4 observed, 59.7 calculated. 

Carbon Disulphide [in steam]. 

m = 0.0495 £=16°.5 A =* 717.8 v = 16.4cc. 

i) 38.6 observed, 38.0 calculated. 

Iodine [in amyl benzoate]. 

m =0.1157 £ = 16°.l /* = 722.3 v = 11.Qqg. 

D 126.4 observed, 127.0 calculated. 

Mercury [heated in lead bath]. 

m — 0.0905 ^ = 16°.0 li = 715.8 v= 1L5-CC. 

D 100.7 observed, 100.0 calculated. 

The student should make several determinations of gas 
density, or if this is not practicable, the instructor should 
make them and let the class make the required calculations. 



CHEMISTRY. 



Elements, Atoms, Classification. 

Elements. — A substance which cannot be resolved into two 
or more different substances is called an element or simple 
body. Gold, silver, iron, copper, lead, tin, sulphur, oxygen, 
nitrogen, and iodine are familiar examples of elementary sub- 
stances. 

Chemical compounds are composed of two or more elements 
held together by their mutual chemical attraction. The 
means at our command for the decomposition of compounds 
are heat, electricity, light, and chemical action. Our experi- 
mental work will make us acquainted with the application of 
these agents not only in isolating elements, but also in effect- 
ing chemical changes. 

We have already learned from the kinetic theory of gases 
that a substance in the gaseous state is composed of very small 
particles, molecules, having definite mass. Most molecules 
are composed of parts which cannot be divided into smaller 
parts. These indivisible particles are called Atoms. Since 
atoms rarely exist in the free state, and only at very high tem- 
peratures, the following definition may be given. An atom is 
the smallest particle of matter existing in combination. 

The evidence of the existence of atoms is chemical, and we 
can therefore best study the atomic theory after having be- 
come familiar with a large number of chemical facts. Since 
we shall use, in beginning the study of chemistry, the terms 



ATOMIC WEIGHTS. 47 

atom and atomic weight, we shall need some idea of their 
meaning. The atoms of an element are alike in all respects; 
the atoms of different elements, that is, different kinds of 
atoms, differ in mass, and more or less in other properties. 
The distinguishing characteristics of a mass of matter are de- 
termined by the properties of the individual atoms which 
compose it. 

An atomic 10 eight is a number expressing the ratio of the 
mass of the smallest particle of an element entering into com- 
bination, to the mass of an atom of hydrogen or half the mole- 
cule of hydrogen. 

It is obvious from the definition that an atomic weight, or 
a multiple thereof, represents the relative weight of an element 
taking part in chemical changes. 

Symbols. — Chemists represent the elements by symbols 
which are the initial letters of their names, or the initial let- 
ter and one other. Some of the symbols are from the Latin or 
other names. A symbol of an element stands for an atom, 
and also for an atomic weight. 

Formulas. — The number of atoms in a molecule of an ele- 
ment is indicated by a number placed below and to the right 
of the symbol: thus, H 3 indicates that the molecule of hydro- 
gen contains two atoms. A compound is represented by the 
symbols of its constituents: thus, hydrogen chloride, a com- 
pound of hydrogen and chlorine, has the symbol HC1. This 
represents one atom and 1 weight of hydrogen, one atom and 
35.5 weights of chlorine. Likewise, the composition of water 
is best shown by the symbol H 2 0, which represents two atoms 
and 2 weights of hydrogen, and one atom and 16 weights of 
oxygen. 

The calculation of the percentage composition of a com- 
pound whose formula- is given is a simple arithmetical pro- 
cess which requires no explanation. 



48 



ATOMIC WEIGHTS. 



The following is an alphabetical list of the 68 elements 
recognized up to the year 1888, together with their symbols 
and atomic weights; the names of the rarer elements are 
in italics. 



Aluminum, 


Al 27 


Molybdenum, . 


Mo 96 


Antimony, 


Sb 120 


Nickel, . 


Ni 58 


Arsenic, . 


As 75 


Niobium, . 


Nb 94 


Barium, . 


Ba 137 


Nitrogen, . 


N 14 


Beryllium, 


Be 9 


Osmium, . 


Os 192 


Bismuth, . 


Bi 208 


Oxygen, . 


O 16 


Boron, 


B 11 


Palladium 


Pd 106 


Bromine, . 


Br 80 


Phosphorus, 


P 31 


Cadmium, 


Od 112 


Platinum 


Pt 195 


Cmium, . 


Cs 133 


Potassium (kalium),. 


K 39.1 


Calcium, . 


Ca 40 


Rhodium, 


Rh 104 


Carbon, 


O 12 


Rubidium, 


Rb 85 5 


Cerium, 


Ce 141 


Ruthenium, 


Ru 103 


Chlorine, . 


CI 35 45 


Samarium, 


Sm 150 


Chromium, 


Cr 52.5 


Scandium, 


Sc 44 


Cobalt, . 


Co 59 


Selenium, 


Se 79 


Copper (cuprum), 


Cu 63.3 


Silicon, . 


Si 28 


Didymium, 


D 145 


Silver (argentum), . 


Ag 107.9 


Erbium, . 


E 166 


Sodium (natrium), . 


Na 23 


Fluorine, . 


F 19 


Strontium, 


Sr 87.5 


Gallium, . 


Ga 70 


Sulphur, 


S 32 


Germanium, 


Ge 72 


Tantalum, 


Ta 183 


Gold (aurum), . 


Au 196.7 


Tellurium, 


Te 125 


Hydrogen, 


H 1 


Thallium, 


Tl 204 


Indium, 


In 113 7 


Thorium, 


Th 232 


Iodine, 


I 127 


Tin (stannum), 


Sn 118 


Iridium, . 


Ir 193 


Titanium, 


Ti 48 


Iron (ferrum), . 


Fe 56 


Tungsten (wolfram), 


W 184 


Lanthanum, 


La 139 


Uranium, 


Ur 239 


Lead (plumbum), 


Pb 207 


Vanadium, 


V 51 3 


Lithium, . 


Li 7 


Ytterbium, 


Yb 173 


Magnesium, 


Mg 24.4 


Yttrium, 


Y 89 


Manganese, 


Mn 55 


Zinc, 


Zn 654 


Mercury (hydrargyn 


lm), Hg 200 


Zirconium, 


Zr 90.7 



CLASSIFICATION. 



49 



Classification of the Elements. — The classification of the 
elements adopted in this book is that known as the Periodic 
System of the Elements. In this system the elements are ar- 
ranged according to increasing atomic weights, and in groups 
whose members are more or less closely related in properties. 
The following is Mend elejeffs table* with a few minor changes: 



Groups. 


Periods. 








I. 


II. 


III. 


IV. 


V. 


VI. 


R 2 

RO 
R 2 3 
R0 2 
R 2 5 
R0 3 
R 2 7 
RO 
to 
R0 4 
R 2 
RO 
R 2 3 
R0 2 


I. 

II. 

III. 

IV. 

V. 

VI. 

VII. 

VIII. 

I. 
II. 
III. 

IV. 

V. 

VI. 

VII. 




Li 7 
Be 9 

Bll 

C12 

N14 
16 
F19 

Na23 
Mg24.4 

A127 

Si 28 

P31 
S32 

CI 35.5 


K39 
Ca40 

Sc44 

Ti48 

V 51 

Cr52.5 
Mn 55 
Fe56 
Co 59 

Ni58 
Cu63.3 
Zn 65.4 
Ga70 
Ge72 
As 75 

Se79 

Br 80 


Rb 85.5 
Sr 87.5 
Y 89 
Zr 90.7 
Nb 94 
Mo 96 

Rul03 
Rhl04 
Pdl06 
Ag 107.9 
Cdll2 
In 113.7 
Snll8 
Sbl20 
Tel25 
1127 


Csl33 
Bal37 
La 139 

Cel41 

Di 145 

Sm 150 














Ybl73 

Tal83 

W184 

0sl92 
Ir 193 
Pt 195 
Aul97 
Hg200 
T1204 
Pb 207 
Bi 20S 




(H 4 C) 
(H 3 N) 
(H a O) 
(HF) 


Th232 


U239 









HI 













Br 166 




(H 4 R) 
(H 3 R) 
(H 2 R) 
(HR) 




R 2 6 
R0 3 
R 2 7 









There are several other forms of tables of the periodic clas- 
sification, but all based on the same principles. The order 
of treatment adopted in this book may be briefly stated. 
First, hydrogen is described; then the seventh group is con- 

* Berichte der deuUsc/wu chemischen Oesellschafi, 1881, p. 3822, 



50 HYDROGEN. 

sidered, together with the compounds its members form with 
each other and with hydrogen. The first, sixth, second, 
fifth, third, fourth, and eighth groups are taken up in the 
order named. As a rule, only those compounds are described 
which an element forms with elements previously studied. 

The student will find it necessary to look up the properties 
of many substances before studying them systematically, and 
in order to find them he should use the index. If the begin- 
ner will in addition frequently refer to larger works, such as 
Roscoe and Schorlemmer's Treatise on Chemistry, and Watt's 
Dictionary of Chemistry, for a description of substances in 
which he is interested, he will early form the invaluable habit 
of using the literature of chemistry. 



Hydrogen, H. 

Atomic Weight 1. Molecule H 2 . 

Hydrogen is a constituent of water, of animal and vegetable 
matter, and enters into the composition of a very large num- 
ber of artificial compounds. Free hydrogen has been detected 
in volcanic gases, but is not found in ordinary atmospheric air. 

Hydrogen is a colorless, odorless gas, and is the lightest 
known form of matter, being 14.44 times lighter than air. 
One liter of hydrogen at 0° and 760 mm. pressure weighs 
0.0896 gram. The density of hydrogen in the alloy it forms 
with the metal palladium is 0.62 ; lithium, the lightest of 
metals, having a density of 0.59. Hydrogen was formerly re- 
garded as a permanent gas, because it was found to remain in 
the gaseous state when subjected to enormous pressure. It 
was first liquefied by Cailletet in December, 1877, and inde- 
pendently by Pictet in the following month. Both experi- 
menters exposed the gas to very low temperatures and great 
pressures. Pictet's result was attained at — 140° and a pres- 
sure of 650 atmospheres. When the apparatus containing 
the liquefied gas was opened steel-gray particles shot out 



HYDROGEN". 



51 



Fig. 4JL 



which rattled on the floor like a metal, the liquid hydrogen 
having been frozen by its own evaporation. The boiling- 
point of liqnid hydrogen is about — 215°. 

Exp. 4. — Weigh accurately a thin glass tube of size shown in Fig. 49. 

Cut the coating from the end of a piece of sodium and push the tube 
into the metal until nearly filled. Cut ' away the 
sodium from about the tube, then push with a 
match-stick the sodium in the tube down about 
2 mm. from the open end. Wipe the tube care- 
fully, weigh quickly, and then place in rock-oil 

in which sodium has been kept. The metal cannot be kept in the air- 
but the small surface exposed at the 

end of the tube will change very 

little during the weighing. 

Seize the tube with forceps and 

place it under the graduated tube A, 

Fig. 50, which has been previously 

filled with water. Hydrogen gas 

will be set free, and in a short time 

the sodium will be dissolved in the 

water. Sink A in the water in B, 

and take the temperature of the 

water, which will be that of the 

gas. Next raise the graduated tube 

until the water in it stands at the 

same level as in the cylinder, thus 

making the pressure of the gas the 

same as that of the air. Note the 

number of cubic centimeters of gas 

and the height of the barometer. 

Calculate * the weight of the hy- 
drogen obtained, and divide the 

weight of sodium taken by it. The 

result will show that 23 weights of 

sodium are required to set free one 

weight of hydrogen from water. 




Fig. 50. 



One experiment was as follows : 

Tube and sodium 0.5214 gram. 

Tube 03335 ' " 

Sodium 0. L879 gram. 

{SM JOOt Of IU.lt />(!;.< .1 



52 HYDKOGEX. 

The atomic weight of sodium (see table of atomic weights) 
is 23; hence the conclusion from the foregoing experiment 
that each atom of sodium sets free one atom of hydrogen in 
the chemical change which occurs when sodium decomposes 
water. The reaction may be formulated as follows : 

N"a + HOH = H + XaOH. 

Sodium, Water, Hydrogen, Sodium hydroxide, 

23 weights. Hydrogen, 2 weights. 1 weight. Sodium, 23 weights. 

Oxygen, 16 " Oxygen, 16 " 

Hydrogen, 1 " 

The equation represents not only that hydrogen is separated 
from water, but also that sodium replaces part of the hydrogen 
in water with formation of sodium hydroxide, a compound 
described later. 

Equations which represent known chemical changes are 
based on the results of experiments in which the weights of 
part or all of the elements reacting have been determined. 
Chemists, however, often express a hypothesis regarding a 
reaction by an equation. 

Exp. 5. — Remove the tube C, Fig. 51, and the rod supporting it, from 
the graduated tube B, and put into G about 0.3 gram of pure zinc, 
accurately "weighed, and best in the form of a thin strip. Replace G 
in B, and while holding the c}iinder D horizontally lay B in it. Next, 
set the apparatus upright, pour some water into D, and fill B with water 
by sucking at A. 

Hydrogen 99.5 cc. at 19°. 

Barometer , . 763 mm. 

Pressure of aqueous vapor at 19° 16.3 

Pressure of hydrogen 746.7 mm. 

746 7 x" 99 5 

MU X 1 + (.00367 x W) = 91A CC - the ™ 1Ume Et °° aDd 76 ° 

mm. pressure. 
91.4 X 0.0000896 (the weight of 1 cc. of hydrogen) = 0.00819 

gram, the weight of the hydrogen obtained. 
0.1879 -f- 0.00819 = 22.94. 
Hence 1 weight of hydrogen was set free by 22.94 weights of sodium. 
In another experiment 22.90 was the result. 



HYDHOGETsT. 



53 



Place in A a drop of solution of platinum chloride and some con 
centrated hydrochloric acid. Allow the acid to run into B, taking 
care that no air enters. After the liquid in B 
has fallen below 0, more acid may be added to 
B. When the zinc has disappeared fill A with 
water and pass the water into B, and repeat the 
washing three or four times in order to remove 
most of the acid from the tube. Fill D with 
water, and raise A until G and the rod fall out. 
Finally, observe the temperature of the water, 
the volume of the gas, and the height of the 
barometer, and from these data calculate the 
relation between the weight of the zinc taken 
and that of the hydrogen obtained. Repeat 
the experiment until the results agree closely, 
and are very nearly 1 of hydrogen to 32.5 
of zinc. 

The apparatus described in the above experi- 
ment is a modification of that devised by 
Keiser,* whose form of apparatus is also good. 



The atomic weight of zinc is 65, and 
the foregoing experiment shows that G5 
weights of zinc set free 2 weights of hy- 
drogen ; that is, 1 atom of zinc replaces 
2 atoms of hydrogen. Hydrochloric acid 
has the formnla HC1, CI standing for 
the element chlorine. The zinc unites 
with the chlorine to form the compound 
zinc chloride, ZnCl 2 , which is dissolved 
in the water. What is the equation 
representing the reaction between zinc 
and hydrochloric acid ? 

Hydrogen is also made by adding sul- 
phuric acid, HJSO.,, diluted with water, 
to zinc, and in other ways. 




a D 



Fin M. 



* American Chemical Journal, vol. vi. 347. 



54 



HYDROGEN. 



Exp. 6.— Place in the bottle A, Fig. 52, about 100 grams of granulated 
zinc, pour through the funnel tube B sufficient water to cover the lower 
end of B, and then add sulphuric acid gradually until gas is evolved 
rapidly. In order to collect the gas over water a jar is filled with water, 
covered with a piece of wet cardboard, and inverted in the water in the 
pneumatic trough D. The jar is then placed over the bole iu the shelf 
of D, and gas is passed by means of the delivery tube under the funnel, 
shown in section by E, when gas will rise and displace the water in 
the jar. 



.. 



^™^,«s^™ 




Fig. 52 



Fill a small jar with gas from the generator A. Remove the jar from 
the trough, holding it mouth downwards, and set fire to the gas. The 
result will probably be an explosion due to the air mixed with the hy- 
drogen which was in the generator. Fill the jar again with hydrogen 
and ignite the gas. If it burns quietly it is sufficiently pure for further 
experiments. 

Exp. 7. — Allow hydrogen gas to escape from the delivery tube G, Fig. 
52, into a dry jar. After a short time remove the jar, set fire to the gas 
at the end of G, and cover the flame with ano h3r dry bottle. Note ob- 
servations. The yellow color of the hydrogen flame in this experiment 
is due to the sodium in the glass. 

Exp. 8. — Mix in a half-pint or pint jar 5 volumes of air and 2 volumes 
of hydrogen, and set fire to the mixture. 

Exp. 9. — Thrust a burning splinter of wood into a jar of hydrogen held 
mouth downwards. 



HYDROGEN. 55 

The burning of hydrogen is one of the chemical properties 
of the element, and the change which occurs is chemical. It 
is better to study the nature of the reaction when we become 
familiar with oxygen, which unites with hydrogen to form 
water. Any substance which yields water when burned with 
air or oxygen contains hydrogen. 

Exp. 10. — Fill two jars with hydrogen, leave one mouth uncovered 
upwards, and the other mouth downwards on the ring of a lamp-stand so 
that its mouth is not closed. After a few minutes apply a light to the 
gas in each bottle. 

Exp. 11. — Fill a jar by holding it over a rapid current of hydrogen 
from the delivery tube G, Fig. 52, and prove that hydrogen has displaced 
the air which was in the jar. 

Exp. 12. — Place a jar containing hydrogen over a similar one filled 
with air, and slowly invert the two. Apply a light to each. If success- 
ful, the gas in the jar which originally contained hydrogen will not 
burn. 

Exp. 13. — Close the bowl of a white clay pipe with a disk of card- 
board and sealing-wax. Connect the stem of the pipe with a narrow 
glass tube by means of rubber tubing. The apparatus is sufficiently tight 
if the tongue is held on the end of the tube after sucking air from the 
pipe. Place the end of the tube in water, and surround the pipe with a 
jar of hydrogen, and after a minute or two remove the jar. Observe 
whether gas bubbles through the water or the water rises in the tube, 
and note explanation of phenomena. Keep the pipe for future experi- 
ments. 

Exp. 14.— Blow soap-bubbles with hydrogen, and hold a flame to some 
of the large bubbles. 

Exp. 15. — When the evolution of hydrogen has ceased in the generator, 
Exp. 6, Fig. 52, filter the solution into a porcelain dish, label it zinc sul- 
phate, and set in the locker. The water will slowly evaporate, and after 
some days crystals will form, having the composition ZnS0 4 -}- 7H a O. 



THE SEVENTH GKOUR 

The elements of this group are the non-metals chlorine, 
bromine, iodine, and fluorine, and the metal manganese. 



Chlorine, CI. 

Atomic Weight, 35.5. Molecule, Cl 2 . 

Chlorine occurs in nature only in combination with other 
elements. Its most abundant compound is common salt or 
sodium chloride. 

It is a greenish-yellow gas about two and a half times heav- 
ier than air, having a density of 35.5 (hydrogen = 1). It 
condenses under cold and pressure to a yellow liquid, boiling 
at -- 33°. 6. Chlorine gas dissolves in about half its bulk of 
cold water, forming a solution known as chlorine water. 

Chlorine is obtained by several different methods. In the 
laboratory it is commonly prepared by warming manganese 
dioxide, Mn0 2 , with strong hydrochloric acid, which react as 
follows : 

Mn0 2 + 4HC1 = 3InCl 2 + 2H 2 + 2C1. 

Chlorine gas is also prepared by heating a mixture of manga- 
nese dioxide, common salt, sulphuric acid, and water, whereby 
chlorine, manganese sulphate, hydrogen sodium sulphate, and 
water result : 

Mn0 2 +2XaCl+3H 2 S0 4 = 2Cl+MnS0 4 +2HXaS0 4 +2H 2 0. 



CHLOBI^E. 



57 



Exp. 16.— The flask A, Fig. 53, holds about a liter. Place in it 100 
grams of pulverized manganese dioxide, 100 grams of common salt in 
coarse crystals, and a cooled mixture of 110 cc of strong sulphuric 
acid and 200 cc. of water. Much heat is evolved on mixing the acid 
and water, and there is danger of acid being thrown out of the vessel 
if water is poured into the acid. Hence the latter should be gradually 
added to the water. The wash-bottle B contains water to absorb hydro- 
chloric acid gas, and C contains concentrated sulphuric acid to dry the 




Fig. 53. 



chlorine. A glass tube from A passes to the bottom of the cylinder D. 
containing a solution of potassium hydroxide to absorb any excess of 
chlorine. The glass stopcock E is to regulate the current of tin 1 gas. 
On gently warming the flask A by means of the water-bath, chlorine 
will come off regularly for a long time When tin 1 gas which pa^srs 
from the bottle C is mostly absorbed by a solution of potassium hydrox- 
ide it is sufficiently free from air for use. 

Fill glass-stoppered jars with chlorine by passing the gas to bottom 
of the jars until the color indicates that the air has been displaced by 



58 



THE SEVENTH GROUP. 



the chlorine. Smear the stoppers with vaseline to make them gas 
tight. As chlorine is irritating and poisonous, experiments with it 
should be conducted under a good hood, or in such a way as to avoid 
inhaling the gas. 

If it is not convenient to collect chlorine under a hood, the apparatus, 
Fig. 54, may be used. The plate contains lime water to absorb any 
chlorine which may escape on tilling the jar under the bell-jar. 




u 



Fig. 54. 



liG 55. 



Exp 17. — Ignite a small jet of hydrogen at the end of the glass tube, 
Fig. 55, tipped with a small tube made of platinum foil. Thrust the 
hydrogen flame into a jar of chlorine standing under the bell-jar (Fig. 
54), and close the mouth of the latter with the rubber stopper A. 

Exp. 18. — Set fire to a large jar of hydrogen, and introduce into it a 
jet of chlorine. The hydrogen will appear to support the combustion 
of the chlorine, while in the preceding experiment the chlorine appeared 
to be the supporter of the combustion. In either case the flame is due 
to the heat evolved by the union of the two gases, forming a compound 
described later. 

Exp. 19. — a. Place in a cylinder, as shown in Fig. 56, a thin test-tube 
about 8 inches long and an inch in diameter, filled with water. Pac-s 
chlorine through the delivery-tube until the test-tube is half filled 
wilh the gas, and then pass in hydrogen until the test-tube is filled 
with the mixed gases, and remove the delivery-tube. The water will 



CHLORINE. 



59 



slowly rise in the test-tube, and nearly fill it after some time, owing to 
the absorption of the product of the slow com- 
bination of the two gases. 

b. Fill the tube again with equal volumes of 
chlorine and hydrogen, cover the top of the 
cylinder with a towel, and expose the mixed 
gases to bright sunshine or the light of burn- 
ing magnesium, when a violent explosion will 
occur. 



The foregoing experiments show that 
chlorine and hydrogen unite at high tem- 
peratures or when exposed to light. The 
compound formed contains both of these 
elements, and is called hydrogen chloride. 

Let us next try to find the relative 
number of volumes of hydrogen and of 
chlorine which unite, and also the volume 
of the hydrogen chloride gas formed. 




Fig. 56. 



20. -—The apparatus, Fig. 57, is filled with a concentrated solu- 
tion of hydrogen chloride in water, known as con- 
centrated hydrochloric acid. The wires from a gal- 
vanic battery are connected with the gas-carbon 
poles. In one tube gas will be evolved rapidly. 
It is hydrogen, and hence will burn. At first but 
little gas will collect in the other tube, and on al- 
lowing it to escape into the air its odor will indicate 
that it is chlorine. After a time the acid will be- 
come saturated with chlorine, and on closing the 
stopcocks equal volumes of hydrogen and chlorine 
will be obtained. 

Exp. 21.— In order to obtain a mixture of equal 
volumes of hydrogen and chlorine, hydrochloric acid 
is subjected to electrolysis in some form o( appa- 
ratus, such as shown by Fig. 58, in which the two 
gases are not separated. The apparatus consists o( 
the tube A, with a rubber stopper, through which 
pass two gas-carbon sticks connected with wires from a battery by 
means of platinum wires. The neck of A passes through a cork 




60 



THE SEVENTH GROUP. 



iii a glass jar which serves as a stand. A is partly filled with 

concentrated hydrochloric acid. A cur- 
rent from 4 to 6 Bunsen cells will cause 
a rapid evolution of the mixed gases, 
and after the acid has become saturated 
with chlorine a mixture of equal vol- 
umes of hydrogen and chlorine will be 
obtained. In order to protect the gases 
from the action of sunlight the appa- 
ratus may be placed in a box, the latter 
being then filled with sawdust to in- 
sure complete exclusion of light. 

Wrap the tube, Fig. 59, in black cloth, 

and then pass through it the gases from the electrolysis of hydrochloric 




Fig. 58. 



Fig. 59. 



I 



acid until the air is expelled, taking pains not to expose the gases to 
daylight. Close the stopcocks, and, without removing 
the cloth, place the tube upright and allow one or two 
cubic centimeters of a solution of potassium iodide to 
flow in through the upper stopcock. Care should 
be taken not to admit air. The chlorine will be ab- 
sorbed by the potassium iodide and iodine liberated. 
Finally, remove the cloth, open one stopcock under 
water in a cylinder, and lower the tube until the re- 
maining gas is under atmospheric pressure The ex- 
periment will demonstrate that half of the volume of 
the mixed gases from the electrolysis of hydrochloric 
acid is chlorine. The hydrogen remaining in the tube 
may be burned. 

Exp. 22. — Pass a mixture of equal volumes of hydro- 
gen and chlorine through a tube, such as shown in Fig. 
60, until the air has been displaced. Then close the upper 
stopcock, disconnect the tube from the generator, and 
quickly close the lower stopcock. Expose the tube to 
bright sunshine or a magnesium light held near the. 
bulb. A slight click will indicate that the gases have Fig. 60. 
combined. Next, open one end of the tube under mercury. Gas 




MOLECULES. 



61 



will not escape nor will mercury enter the tube, showing that the 
pressure of the gas in the tube is the same as before combination. 
Finally, fill one end of the tube beyond the stopcock with water, 
place under water, and open the lower stopcock. Water will absorb 
the gas quickly and fill the tube. The experiment shows that equal 
volumes of hydrogen and chlorine unite to form a compound (hydrogen 
chloride) which is soluble in water. 

Molecules. — The combination of hydrogen and chlorine may- 
be represented thus : 



Hydrogen. 



+ 



1 Vol. 



Chlorine. 



Hydrogen Chloride. 



1 Vol. 



2 Vols. 



According to Avogadro's law, the two volumes of hydrogen 
chloride contain twice as many gas molecules as one volume 
of hydrogen. Since each molecule of the new gas contains 
both hydrogen and chlorine, it follows that in the reaction 
the molecules of both of these substances must have separated 
into two parts, each of the parts of the hydrogen molecule 
uniting with one of the parts of the chlorine molecule. Thus, 
since the experiment proves that the molecules of both sub- 
stances are divisible, each kind of molecule must contain 
more than one atom. The simplest supposition is that each 
is composed of two atoms, but the combination can be equally 
well explained by supposing the hydrogen molecule to con- 
tain four or more atoms and the chlorine molecule an equal 
number. A consideration of the hydrogen chloride molecule 
and a comparison of it with the hydrogen molecule will be o( 
service. Chlorine gas is 35.5 times heavier than hydrogen, and 
one volume of hvdroeren weighing' 1 unites with one volume of 



chlorine weievhini 



the atomic weights o( the two are re- 



62 THE SEVEXTH GKOUP. 

spectively 1 and 35.5 : hence the hydrogen chloride molecule 
consists of an equal number of atoms of each element. It may 
be represented by HC1, H 2 C1 2 , H 3 C1 3 , or a larger number of 
atoms. If it contain two or more atoms of hydrogen we should 
expect to find some chemical changes in. which only part of 
the hydrogen is replaced. On the contrary, all the hydrogen 
of hydrogen chloride is replaced where the latter is acted upon 
by metals. From this chemical view we conclude that the 
hydrogen chloride molecule contains but one atom of hydro- 
gen and one atom of chlorine, and that it is represented by 
the formula HC1. If the hydrogen chloride molecule is HC1, 
the hydrogen molecule is HH or H 2 , as is evident from the 
following reasoning. Let us suppose one volume of hydrogen 
to contain N molecules. A like volume of chlorine will also 
contain, according to Avogadro's law, X molecules, and the 
two volumes of hydrogen chloride resulting from the combi- 
nation of the two gases will contain 2X molecules of hydro- 
gen chloride. Each molecule of hydrogen chloride contains 
one atom of hydrogen, and the 2N molecules contain 2X 
atoms of hydrogen which were originally contained in X 
molecules of hydrogen. Hence each hydrogen molecule con- 
tains two atoms and is represented by H 2 . The same argu- 
ment leads to the conclusion that Cl 2 represents the chlorine 
molecule. 

The ratio of the weight of any gas molecule to that of the 
hydrogen molecule is the ratio of the density of the gas to 
that of hydrogen at the same pressure and temperature, as 
has been shown, p. 38. Since the smallest quantity of matter 
which is known to act as an independent body in chemistry is 
the atom of hydrogen, i.e. one half the hydrogen molecule, 
it is convenient to choose that as the unit by which molecular 
weights are measured ; hence we have the rule — 

The molecular weight of any substance is equal to twice its 
gas density. 

Recalling again the union of hydrogen and chlorine : 



MOLECULAR WEIGHTS. 63 

Hydrogen. Chlorine. Hydrogen Chloride. 




1 Vol. 



1 Vol. 



18.25 
Weights. 



18.25 
Weights. 



Vols. 



The gas density of hydrogen chloride has been found by 
experiment to be 18.25. The molecular weight of hydrogen 
chloride is, therefore, 36.5. This corresponds to the weight 
of one atom of chlorine, plus the weight of one atom of hy- 
drogen ; thus, 35.5 -f- 1 = 36.5. 

The molecular weight of chlorine, according to the rule, 
equals 35.5, the gas density of chlorine, multiplied by 2 or 71, 
which is the weight of two atoms of chlorine. The molecule 
of chlorine is therefore Cl 2 . The following rule requires no 
further explanation : 

The number of atoms in a molecule of an element is found 
by dividing its molecular weight by its atomic weight. 

Thus far we have taken for granted the atomic weight of 
chlorine. If we accept the conclusion arrived at from the 
consideration of the union of hydrogen and chlorine that the 
chlorine molecule contains two atoms, then the atomic weight 
of chlorine equals 35.5, or the molecular weight derived from 
the gas density of chlorine divided by 2. When we study 
the subject of atomic weights we shall have more complete 
reasons for regarding the atomic weight of chlorine as 35.5. 

Molecular Weights of -Compounds. — The molecular weight 
of a compound is twice its gas density, according to the rule 
previously stated. Since many compounds do not gasify 
without decomposing, or at temperatures too high to allow a 
determination of gas density, their molecular weights arc un- 
known. In such cases it is customary to assign to a com- 
pound a formula representing the least mass which is sup- 



64 THE SEVENTH GROUP. 

posed to represent all the properties of the compound. Such 
formulas are often for convenience assumed to represent 
molecules. 

Analysis, Synthesis. — A chemical analysis is the separation 
of a compound into its components ; as, for example, the 
separation of hydrochloric acid into hydrogen and chlorine. 
Qualitative analysis determines the constituents of a sub- 
stance, and the quantities or relative proportions of the 
several constituents are found by quantitative analysis. The 
synthesis of a compound is the uniting of the elements which 
compose it; as, for example, the combining of hydrogen and 
chlorine to form hydrochloric acid. 



Chlorine unites with metals forming compounds termed 
chlorides. It combines energetically with hydrogen, remov- 
ing the latter from many compounds with formation of hydro- 
chloric acid. 

Moist chlorine is a powerful bleaching agent, but the dry 
gas has ordinarily no effect on coloring matters. The bleach- 
ing action of chlorine is due, wholly or in part, to the with- 
drawal of hydrogen from water, setting free oxygen which 
oxidizes or destroys the colored compounds. Oxygen in the 
ordinary state does not act as a bleaching agent. If, however, 
it is in the nascent state, i.e., when the atoms of this element 
are just set free from compounds, it is more active chemically, 
and will bleach coloring matter. Later we shall learn that the 
molecules of ordinary oxygen contain two atoms. The atoms 
of ox} r gen in the nascent state are supposed to be free, having 
a marked tendency to combine with other substances present, 
or with each other. 

The following equation represents the action of chlorine on 



CHLORIDE. 65 

water, one atom of oxygen being set free from one molecule of 
water: 

H 2 + Cl 2 = 2HC1 + 0. 

Chlorine gas is a good disinfectant, destroying germs and 
the noxious compounds of decay. 

Exp. 23. — a. Place some Dutch leaf (an alloy of copper and zinc beaten 
into thin sheets) in a pint jar of chlorine. 

b. Drop some pulverized metallic antimony into chlorine. 

Exp. 24. — Place a dry piece of phosphorus on a deflagrating spoon 
made of chalk and copper wire, and thrust into chlorine. 

Exp. 25. — Fill a large thin glass tube, closed at one end, with a satu- 
rated solution of chlorine in water, and place the open end of the tube 
in a jar containing some chlorine water, taking care that no air enters 
the tube. Set the apparatus in bright sunshine. After a time minute 
bubbles of gas will be seen, and in course of a day sufficient gas will 
collect in the tube to test. The gas thus obtained is oxygen, and will 
cause a glowing splinter of wood to inflame and burn when thrust 
into it. 

Exp. 26. — Saturate a piece of thin filter- paper held by a wire with 
boiling turpentine, and quickly plunge into a jar of chlorine. The tur- 
pentine wiU inflame. It is a compound of carbon and hydrogen. The 
latter unites with the chlorine, and the carbon is separated as soot. 

Exp. 27. — Leave a piece of dry pink calico in a stoppered jar of dry 
chlorine for an hour or longer. Then add a little water. Note the re- 
sults. 

Exp. 28 — Place in chlorine water pieces of calico, printed paper, arid 
paper with lead-pencil and writing-ink marks. Printing ink contains 
lamp-black, a form of carbon, and the lead of a lead-pencil is chiefly 
graphite, another modification of carbon. Free carbon is not attacked 
by chlorine. 

Exp. 29. — To a few drops of a solution of silver nitrate in a test-tube 
add some chlorine water. The white precipitate obtained is silver 
chloride, AgCl. 

The term jyrecipitaie, which means literally something thrown down, 
is applied to solids which separate when two solutions are mixed. 

Exp. 30.— Dissolve a little common salt (sodium chloride, NaCl) in 
water and add to the solution silver nitrate. Also add silver nitrate to 
very dilute hydrochloric acid, in both eases silver chloride will be 
5 



6Q THE SEVEXTH GROUP. 

precipitated. The experiments show that silver nitrate may be used to 
detect chlorine, but later we shall learn that a few other substances, 
less common than chlorine, also yield white precipitates with silver 
nitrate. 

Hydrochloric Acid or Hydrogen Chloride, HC1, is a color- 
less gas, Terr soluble in water. The solution is commonly 
known as hydrochloric or muriatic acid. One volume of water 
at 0° absorbs 503 volumes of the gas at 760 mm. pressure. 
At ordinary temperatures about 450 volumes are absorbed, 
forming a solution having a density of about 1.2, and contain- 
ing approximately -±0 per cent of hydrogen chloride. A con- 
centrated solution of the acid fumes strongly in the air, and 
on heating gives off hydrochloric acid gas, the temperature of 
the liquid rising to 110°, when an aqueous acid containing 
20.24 per cent of HC1 distils unchanged at 760 mm. pressure. 
A more dilute acid loses water on boiling until it attains the 
same strength. The observed density of hydrochloric acid gas 
is 18.25 (hydrogen = 1). 

Hydrochloric acid is made in enormous quantities by heating 
a mixture of common salt and sulphuric acid, and absorbing 
the gas by water. 

In the reaction hydrochloric acid and hydrogen sodium 
sulphate are formed, thus: 

(1) XaCl + H 2 S0 4 = HC1 + HXaS0 4 . . 

If sufficient salt is taken and the mixture more strono-lv 
heated, the hydrogen sodium sulphate reacts with the salt as 
follows: 

(2) HXaS0 4 + XaCl = HC1 + Xa 2 S0 4 . 

The hydrogen sodium sulphate of the first reaction is more 
readily soluble in water than the sodium sulphate obtained by 
the second reaction. 



HYDROCHLORIC ACID. 



67 



How many grams of hydrogen chloride can be made from 
50 grams of sodium chloride, and how much sulphuric acid 
will be required if the salt and acid are used in the propor- 
tion required by the first equation ? 

Exp. 31. — Place 100 grains of coarse crystals of common salt in a flask 
having a capacity of a liter, and pour in a cool mixture of 100 cc, of 
concentrated sulphuric acid and 25 cc. of water. Connect the flask with 




Fig. 61. 



three Woulfe bottles, Fig. 61, containing water. Heal the flask gently 
with a lamp as long as gas is freely evolved. Test the liquid in each 
bottle with blue litmus paper. 

Exp. 32.— The apparatus shown in Fig. 62 may be used to illustrate 
the rapid absorption of hydrochloric acid gas by water. The strong 
glass globe A is filled by passing a rapid current of hydrochloric acid 
gas into it for 15 or 20 minutes. The bottle B contains water colored 



68 



THE SEVENTH GROUP. 



blue with a solution of litmus. Open the pinchcock D, and blow 
into the tube C for an instant, to force the water into A. 

Exp. 33.— Fill the cylinder A, Fig. 
63, with gas obtained by heating fum- 
ing hydrochloric acid contained in the 
flask B. When acid fumes escape free- 
ly from the cylinder remove the de- 
livery-tube, and quickly place the mouth 
of the cylinder under water. The hole 
in the cork G is about twice the diame- 
ter of the delivery-tube passing through 
it. 

The composition and physical 
properties of hydrochloric acid 
have been considered, while as yet 
but few of its chemical properties 
have been noticed. Its acidic 
character and deportment towards 
a class of compounds known as 
;.. bases can best be made evident by 
f the following experiments : 

Exp. 34. — Dilute a few drops of hy- 

Fig 62 

drochloric acid with a test-tube full of 
water. Taste a drop of the very dilute acid, and also try the action of 
it on blue litmus paper. 

Exp. 35.— Take a stick of sodium hydroxide about two inches long, 
notice its action on the fingers, and then dissolve it in a test-tube half full 
of water without applying a lamp flame. Cool the solution by placing the 
test-tube in cold water. Dilute a few drops of the solution largely with 
water, taste, and test with litmus paper reddened with very dilute acid. 
Pour one third of the original solution of sodium hydroxide into another 
test-tube, and add to it some hydrochloric acid. The rise in temperature 
indicates that a change has occurred. Transfer the solution to a porce- 
lain dish, and add to it most of the original solution of sodium hydroxide. 
Then add hydrochloric acid to the contents of the dish until the solution 
does not change the color of a red or a blue litmus paper. When this 
is attained, the acidic properties of the hydrochloric acid and the basic 




HYDROCHLORIC ACID. 69 

properties of the sodium hydroxide used have disappeared. The acid 
and base have mutually neutralized each other. The taste of the solu- 
tion shows that it contains common salt. Filter the solution, and evap- 
orate slowly over a lamp until a dry residue remains. If a solution of 




Fig. 63. 



the salt is allowed to evaporate at ordinary temperature crystalline cubes 
will be obtained. 

The chemical change in which common salt, NaCl, is 
formed from sodium hydroxide, NaOH, and hydrochloric 
acid is represented by the equation — 

NaOH + HOI = NaOl + 11,0. 

ILO is a molecule of water formed in the reaction. The water 



70 THE SEVEXTH GROUP. 

which held the acid and base in solution served as a medium 
in which the changes occurred, but did not take part in them, 
and hence is not represented in the equation. The hydrogen 
of the hydrochloric acid has been exchanged for sodium, 
or, as more commonly expressed, has been replaced by sodium. 
When metallic zinc dissolves in hydrochloric acid the hydro- 
gen replaced escapes as a gas. If a compound of zinc and 
oxygen known as zinc oxide, ZnO, is dissolved in hydrochloric 
acid the hydrogen of the acid is replaced as before by the 
zinc, but instead of escaping, unites with the oxygen of the 
zinc oxide to form water, thus: 

ZnO + 2HC1 = ZnCl 2 + H 2 0. 



Bromine, Br. 

Atomic Weight, 80. Molecule, Br 



Bromine is not found in the free state in nature, but al- 
ways in combination with metals. It exists in small quanti- 
ties in many salt springs and in sea water. A considerable 
part of the bromine in commerce is obtained from salt springs 
in West Virginia. Bromine is a dark-red volatile liquid, 
freezing at — 24°. 5 and boiling at 63°. The density of liquid 
bromine is 3.18 at 0°. The gas density of bromine at 100° is 
80, while at about 1500° it has been found to be 57 to 67, 
showing that diatomic bromine molecules dissociate at high 
temperatures. The calculated gas density of a mixture of 
equal molecules of Br 2 and Br, is GO. 

Bromine is very soluble in chloroform, and is soluble in 30 
parts of water. 

Exp. 36. — Drop a thin glass bulb containing a few cubic centimeters of 
bromine into a tall cylinder, and cover the cylinder with a ground-glass 
plate. The vapor of bromine, though over five times heavier than air, 



IODINE. 71 

will soon rise to the top of the cylinder. In handling bromine care 
should be taken to avoid inhaling its vapor, as it is poisonous, and very 
irritating. Liquid bromine is corrosive to the skin. 

Exp. 37.— Shake up bromine and water in a stoppered bottle to make 
bromine water. 

Bromine may be obtained from bromides by the action of 
manganese dioxide and sulphuric acid in the same way that 
chlorine is obtained from sodium chloride. In chemical de- 
portment bromine resembles chlorine, uniting with metals to 
form bromides analogous to the chlorides. It combines with 
hydrogen to form hydrobromic acid, but the union does not 
take place in sunlight, as is the case with chlorine and hydro- 
gen. If hydrogen mixed with bromine vapor is burned in the 
air a small quantity of hydrobromic acid will be formed. 

Hydrobromic Acid or Hydrogen Bromide, HBr, is a colorless 
gas, very soluble in water. It cannot well be made by acting 
on bromides with sulphuric acid, owing to its decomposition by 
sulphuric acid with separation of bromine. It is best pre- 
pared by adding bromine to red phosphorus and w r ater, when 
the following reaction occurs: 

P + 5Br + 4H 2 = 5HBr + H 3 P0 4 . 

The gas is passed through a tube containing red phos- 
phorus to free it from bromine vapor. 



Iodine, I. 

Atomic Weight, 127. Molecule, I a . 

Iodine occurs sparingly in nature. It is found in salt bods 
and springs in minute quantities. It cannot be profitably ex- 
tracted from sea water, but certain sea weeds take it from the 
water, and from their ashes or kelp iodine is obtained. The 



72 THE SEVENTH GROUP. 

mother-liquors of Chili nitre, NaX0 3 , are also a source of 
iodine. 

Iodine is an almost black crystalline solid, with a metallic 
lustre,, having a density of 4.95 at 17°. It melts between 113° 
and 115°, and boils above 200°. Its vapor when pure appears 
deep blue by transmitted light, but when mixed with air it 
has a reddish-violet color. Iodine volatilizes slowly at ordinary 
temperature, and its odor is similar to but feebler than that 
of chlorine and bromine. Iodine is soluble in chloroform, in 
a solution of potassium iodide, and in alcohol, but is almost 
insoluble in water. 

The gas density of iodine below 500° is 127, corresponding 
to the molecule I 2 . With increasing temperature the density 
diminishes, and above 1500° it has been found to be 65.7 ; the 
calculated density of molecules of I x is 63.5. The diatomic 
molecules of iodine dissociate at a sufficiently high tempera- 
ture into monatomic molecules. 

Iodine forms compounds analogous in composition and prop- 
erties to chlorine and bromine compounds. It unites with 
metals to form iodides. Chlorine and bromine displace iodine 
from metallic iodides ; for example — 

KI + CI = KC1 + I. 

The characteristic reaction of iodine is the blue color it im- 
parts to starch-paste. Combined iodine, however, does not 
blue starch. 

Exp. 38. — Place a few crystals of iodine in a dry test-tube, and warm 
in lamp flame until iodine vapor fills the tube ; observe the color, and 
allow to cool. Small crystals will be formed on the sides of the tube. 
Next add water, and heat to boiling. Notice whether iodine escapes with 
the steam. Preserve the aqueous solution for further experiments. 

Hydriodic Acid or Hydrogen Iodide, HI, is a colorless gas 
similar in properties to hydrogen chloride and hydrogen bro- 
mide. Iodine and hydrogen unite when passed over hot 



HYDRIODIC ACID. 



73 



platinum sponge, but practically hydrogen iodide is obtained 
by the reaction which occurs between phosphorus, iodine, 
and water : 

P + 51 + 4H 2 = 5HI + H 3 P0 4 . 

Hydrogen iodide is very soluble in water, the solution 
saturated at 0° having a density of 1.99. The aqueous solu- 
tion is colorless, but soon reddens on exposure to air and light, 
owing to separation of iodine. Hydrogen iodide is readily 
decomposed by heat. The influence of a small amount of 
matter on a compound is well illustrated by hydriodic acid 
containing in 128 parts only 1 part of hydrogen. Hydrogen 
iodide gas is decomposed by mercury, hence it can only be 
collected by displacing air. 

Exp. 39. — Place in the retort, Fig. 64, a few grams of red phosphorus, 
and drop in slowly by means of the tap funnel a solution of 2 parts of 




Fig. 64. 



iodine in 1 part of hydriodic acid of density 1.7. Hydrogen iodide will 
be evolved without application of heat at first; later, gentle warming 
will be necessary. To free the gas from iodine vapor, moist red phos- 
phorus is placed with a loose filling of asbestus in the neck of the 
retort. Fill two or three glass-stoppered jars with the gas; then pass 
the gas into water contained in a test-tube standing in a jar of cold 
water. Preserve the solution for future use. 



74 THE SEVENTH GROTIF, 

Exp. 40. — a. Pour hydrogen iodide gas from a jar on a large Bunsen 
flame. The gas will be decomposed by the heat, and a brilliant cloud of 
iodine will be seen. b. Allow a colorless solution of hydrogen iodide to 
stand in light to find if decomposition occurs. 

Iodine Monochloride, IC1, and Iodine Trichloride, IC1 3 . — 

"When dry chlorine is passed into a jar containing iodine, a 
dark-brown liquid results, from which IC1 crystallizes on cool- 
ing. With an excess of chlorine, lemon-yellow crystals of 
IC1 3 form, which give off chlorine at 25°, leaving the mono- 
chloride. 

Exp. 41. — To a granule of potassium bromide in a test-tube add con- 
siderable chlorine water. The change in color will indicate that bro- 
mine is set free : 

KBr -f CI = KC1 + Br. 

Add to the solution a few drops of chloroform, and shake violently. 
Note observations. 

Exp. 42. — Add chlorine water gradually to a small granule of potas- 
sium iodide. The first portions of iodine set free will dissolve in the 
solution of potassium iodide, and when sufficient chlorine water has 
been added a black precipitate of iodine will separate, since iodine is 
nearly insoluble in a solution of potassium chloride. 

Exp. 43. — Dissolve a few crystals of iodine in a little alcohol, and 
place some of the solution on the hand. The stain may be removed by 
ammonia water. 

Exp. 44. — Place in half a test-tube of water not more starch than can 
be taken up on the end of a penknife blade ; shake thoroughly, and 
heat to boiling, with frequent shaking to avoid burning the starch on the 
bottom of the tube. The starch granules burst in warm water, and a 
thin, translucent starch-paste will result. 

Add a few drops of starch-paste to chlorine water, bromine water, 
and iodine water in separate test-tubes, and note any changes. To the 
tube containing iodine add an excess of chlorine water. 

Exp. 45. — Dilute a few drops of solution of potassium iodide* with 

* Experience has shown that a solution of potassium iodide which has stood for 
a time in a laboratory is liable to contain free iodine. The instructor should test 
the solution beforehand, and, if it reacts for iodine, he may cautiously add potas- 
sium hydroxide until no reaction for iodine appears on testing with starch-paste. 



FLUOBIKE. 75 

water, and add a little starch-paste. Note change, if any; then add a 
drop of chlorine water, and finally sufficient chlorine water to remove 
color. 

Exp. 46. — To a solution of potassium bromide and starch-paste add 
chlorine water. 

Exp. 47. — To a very dilute solution of potassium iodide containing 
starch-paste add a little and then an excess of bromine water. 

Represent by equations the reactions in the foregoing experiments in 
which chlorine liberates bromine and iodine, and in which bromine 
liberates iodine. Chlorine, iodine and water react to form iodic acid, 
HI0 3 : 

I + 5C1.+ 3H 2 = HI0 3 + 5HC1. 

This explains why an excess of chlorine discharges the blue color of 
iodine and starch, as iodic acid does not color starch blue. 



Fluorine, F. 

Atomic Weight, 19. Gas Molecule unhwiun. 

Fluorine occurs combined with calcium as fluor spar or 
nuorite, CaF 2 , which is abundant in some localities. It is 
also a constituent of cryolite, 3NaF -j- A1F 3 , found in Green- 
land, and of several other minerals. Sea water and the waters 
of many mineral springs contain it, and it has been found in 
the teeth and other parts of the body. 

Fluorine resisted many attempts to isolate it until the year 
1886, when Moissan succeeded in electrolyzing anhydrous 
hydrogen fluoride kept at a low temperature, and obtained 
free fluorine as a colorless gas. It combines explosively with 
hydrogen in the cold, and sets free chlorine from cold potas- 
sium chloride. Crystallized boron and silicon, arsenic, anti- 
mony, sulphur, and iodine burn brilliantly in fluorine gas. 
The metals also burn in it, but loss violently. Cork chars 
and ignites in the gas, and alcohol, benzene, and turpentine 
are inflamed by it. It decomposes water with formation of 
hydrogen fluoride and ozone. 



76 THE SEVENTH GROUP. 

, Hydrofluoric Acid or Hydrogen Fluoride, HF and H 2 F 2 .— 

Hydrogen fluoride is a colorless liquid, boiling at 19°. 5, and 
solidifying at — 102°. 5 to a transparent crystalline mass, 
which melts at —92°. 3. Its gas density at 30°. 5 has been 
found to be 19.66, showing that at this temperature the 
molecule is H 2 F 2 . At higher temperatures the density of the 
gas is about 10, corresponding to the molecule HF. Hydrogen 
fluoride mixes in all proportions with water. Care is neces- 
sary in experimenting with hydrofluoric acid, as it is very 
poisonous when inhaled, and when pare or in very concen- 
trated solution blisters the skin, producing sores which heal 
slowly. A dilute aqueous solution does not attack the skin. 
Anhydrous hydrofluoric acid is without action on glass, but 
if a trace of water is present it corrodes glass. The aqueous 
acid dissolves glass, thus differing in this respect from all 
other acids. The acid does not attack gold, platinum, or 
silver, and the concentrated aqueous solution of it may be 
kept for several years in a leaden jug, which, however, will 
slowly be corroded. Gutta-percha bottles are used for hold- 
ing it, but in time they crack. The dilute hydrofluoric acid 
of commerce is transported in wooden casks. 

Hydrofluoric acid is used in the laboratory as a solvent, 
and in the arts both the gas and its solution are used 
for etching glassware. The anhydrous acid is best pre- 
pared by heating hydrogen potassium fluoride, HFKF, in a 
platinum retort, and condensing the vapors in a platinum 
vessel immersed in a freezing mixture. The common method 
of making hydrofluoric acid is by heating fluor spar, CaF 2 , 
with concentrated sulphuric acid, when hydrofluoric acid and 
calcium sulphate are produced : 

CaF 2 + H 2 S0 4 = CaS0 4 + 2HF. 

Nearly anhydrous acid distils, which may be collected in a 
cooled receiver, or conducted into water if an aqueous solution 
is desired. 



SUMMAEY OF THE HALOGENS. 77 

Exp. 48. — Cover the convex side of a small watch-glass with a thin 
coating of melted beeswax, and, when cool, mark lines through the 
wax, thus exposing the glass. Place a thimbleful of powdered fluor 
spar in a leaden cup, add sufficient sulphuric acid to make a thin paste, 
and heat cautiously with a small flame until fumes appear on holding 
the stopper from the ammonia bottle over the dish. Place a cardboard 
ring on the cup to keep the warm metal from contact with the wax, and 
on the ring put the watch-glass, with the wax side exposed to the fumes. 
Pour a little ammonia water into the watch-glass to keep the wax from 
melting, and also to combine with any acid which may escape. After 
half an hour remove the wax from the glass by warming, and then rub- 
bing with a piece of paper. 

Exp. 49.— Coat a tube with wax, and graduate it to millimeters. Then 
cover the graduation with a strip of thin filter paper, and moisten the 
paper with strong, but not fuming, hydrofluoric acid by means of a 
camel's-hair brush. In a short time the etching will be completed. 



Summary of the Halogens. 

Fluorine,^ chlorine, bromine, and iodine are known as the 
halogens (salt producers), as they all form with metals com- 
pounds analogous to sea salt, NaCl. They constitute, together 
with manganese, the seventh group of the periodic system. 
Manganese is a metal like iron in many respects, and possess- 
ing few characteristics common to the group. The halogens 
are a well-defined group of elements, exhibiting close analogy 
in their chemical deportment. They unite with hydrogen to 
form the compounds HF, HOI, HBr, and HI, all of which arc 
strong acids. The sodium and potassium halides are N'aF. 
NaOl, NaBr, Nal, KF, KOI, KBr, and KI. The halogensin 
their analogous compounds present a gradation in the energy 
with which they enter into combination, corresponding to the 
differences of their atomic weights. Fluorine, with the lowesi 
atomic weight, combines with hydrogen at common tempera- 
ture; chlorine under the influence of light or heat : bromine, 
less readily; and iodine, with the highest atomic weight 



78 



THE SEVENTH GROUP. 



scarcely combines directly with hydrogen. Fluorine displaces 
chlorine from potassium chloride, and presumably bromine 
and iodine from potassium bromide and iodide. Chlorine 
displaces bromine and iodine, and bromine displaces iodine. 
Fluorine decomposes water rapidly, chlorine slowly, under the 
influence of light, and bromine has little and iodine no action 
upon water. Hydrogen iodide in aqueous solution decomposes 
readily, hydrogen chloride in concentrated solution is slightly 
decomposed by long exposure to light, while hydrogen fluoride 
is unchanged. The following table shows a gradation of physi- 
cal properties with increasing atomic weights : 

Atomic Gas Color. Melting Boiling 

Weight. Density. point. point. 

Fluorine 19 Colorless. .... 

Chlorine 35.5 35.5 Greenish yellow - 33°. 6 

Bromine 80 80 Red. -24.5 63° 

Iodine 127 127 Black. 113 above 200° 



Manganese, Mn. 

Atomic Weight, 55. Density, 8. 

Pyrolusite, or black oxide of manganese, Mn0 2 , and man- 
ganite, Mn 2 3 -j- H 2 0, are the chief ores of manganese. It 
occurs in many minerals, and has been found in small quan- 
tities in animals and plants. 

Metallic manganese is obtained by reducing the oxides by 
means of carbon at a white heat. The pure metal is hard, 
brittle, and has a grayish-white lustre. It oxidizes in air, 
decomposes hot water, and dissolves in acids. Impure man- 
ganese, containing iron, carbon, and silicon, is not affected by 
air. 



Manganous Chloride or Manganese Dichloride, MnCl 2 , re- 
sults from the direct union of its elements. A solution of the 



MANGANESE. 79 

salt is prepared by dissolving any of the oxides of manganese 
in hot concentrated hydrochloric acid. The solution on cool- 
ing or spontaneous evaporation deposits pink crystals of 
MnCl 2 + 4H 2 0, which lose two molecules of water in dry air, 
and deliquesce in moist air. The salt is very soluble in water, 
and readily soluble in alcohol. 

The compounds MnBr 2 , Mnl 2 , and MnF 2 are easily ob- 
tained. 

Manganese Tetrachloride, MnCl 4 , has not been isolated. 
Manganese dioxide, Mn0 2 , dissolves in cold concentrated 
hydrochloric acid. It is supposed that the tetrachloride is 
formed as follows : 

Mn0 2 + 4HC1 = MnCl 4 + 2H 2 0. 

The dark-brown solution soon decomposes, and for each 
molecule of Mn0 2 two atoms of chlorine are set free, man- 
ganese dichloride remaining in the solution. 

Manganese Tetrachloride, MnF 4 , has been obtained in com- 
bination with potassium fluoride as 2KF.MnF.. 



VALENCE. 

We have seen that there are good reasons for supposing the 
molecule of hydrogen chloride to be composed of one atom of 
hydrogen and one of chlorine, and represented by the formula 
HC1. Later we shall learn that there are equally good reasons 
for regarding the molecule of water as H 2 0, that of ammonia 
as H 3 X, and that of methane as H 4 C. These four compounds 
show that chlorine, oxygen, nitrogen, and carbon differ in 
their power to hold in combination atoms of hydrogen. 
Chlorine unites with one, oxygen with two, nitrogen with 
three, and carbon with four atoms of hydrogen. The four 
compounds named illustrate the general fact that the atoms 
of every element possess the power of holding in combination 
one or more atoms. This property is termed valence, and may 
be defined thus : 

The valence of an atom is its capacity to hold in combination 
other atoms or groups of atoms (compound radicals). 

Hydrogen never appears to possess a valence greater than 
one, and is therefore taken as a measure of valence. Most 
elements do not combine with hydrogen alone, and their 
valence is measured in their combinations with elements whose 
valence is measured by hydrogen. Silver, for example, does 
not unite with hydrogen, but unites with chlorine in the pro- 
portion of one atom to one atom, and is therefore regarded as 
having a valence equal to that possessed by an atom of hydro- 
gen. 

Valence is expressed by calling hydrogen and chlorine uni- 
valent or monads, oxygen bivalent or a dyad, nitrogen tri- 
valent or a triad, and carbon tetravalent or a tetrad. For higher 



VALENCE. 81 

valences we have the terms pentavalent or pentad, hexvalent 

or hexad, and heptavalent or heptad. Valence is indicated 

by Koman numerals or bars attached to symbols, thus: 

ii 

or -0-. Water is represented by H-O-H, ammonia by 



H H\ 

H-^N", methane by §^>C, or 

M H/ H 



H\ H 

XT' A I 

These are termed 



graphic formulas. The bars represent what is often vaguely 
called bonds or units of affinity. Thus we say that chlorine 
possesses one, oxygen two, nitrogen three, and carbon four 
bonds. The graphic formulas above are not intended to in- 
dicate the positions of the atoms in the molecules, but simply 
their valence, and the relations the atoms bear to each other. 
In the case of chlorine, oxygen, nitrogen, and carbon their 
valence is measured by the number of atoms of hydrogen which 
combine with each. When we consider molecules composed 
of two elements, each of which exhibits towards hydrogen a 
valence greater than one, we find, according to our definition, 
that an element may have a valence differing from that shown 
in its compound with hydrogen. For example, the molecule 
of carbon dioxide contains one atom of carbon and two atoms 
of oxygen. Assuming that both atoms of oxygen bear the 
same relation to the atom of carbon, we may write the graphic 
formula thus: O-C-O, in which carbon is a dyad. Philo- 
sophically there is as much reason to regard carbon as a dyad 
in carbon dioxide as there is to consider it a tetrad in methane. 
There is, however, this practical objection to viewing carbon 
as a dyad in carbon dioxide — the formula O-C-0 does not 
exhibit the same valence of oxygen and carbon which appears 
in the hydrogen compounds. But by indicating the valence 
each clement possesses, measured by hydrogen, by the formula 
= — 0, we have the same valence for carbon in carbon 
6 



82 VALENCE. 

dioxide as in methane, and also represent the atom of oxygen 
as equivalent to two atoms of hydrogen. 

The halogens towards hydrogen and the metals are mo- 
nads, as in H-Cl, H-Br, H-I, H-F, Xa-Cl, K-Ol, and 
Mn</-n. Iodine, however, exhibits a higher valence in iodine 

/ C1 

trichloride, I^-Cl This is an illustration of one of the chief 
\C1. 

difficulties in the hypothesis of valence, namely: that the 
valence of an element may vary, or is only shown in certain 
compounds. Some authorities regard true valence as the great- 
est capacity of an atom to hold in combination other atoms. 
Since a chemical compound results from the mutual relations 
of the elements composing it, we may say that the valence ex- 
hibited also depends upon the mutual relations of its elements. 
According to this view, an element may exhibit a variable val- 
ence, as illustrated in case of iodine. Other difficulties in the 
imperfect hypothesis of valence cannot be profitably discussed 
now. The hypothesis has been, and is, of great service in 
classifying and formulating compounds. It is also a help in 
studying chemistry, and therefore it has been taken up before 
the student has become familiar with a large number of com- 
pounds, a knowledge of which is necessary in order to under- 
stand a more complete discussion of the subject. We shall 
hereafter in this book use the hypothesis of valence, and use 
graphic or constitutional formulas for many compounds, more 
especially for the more common ones. 



THE FIKST GEOUP. 

The members of this group are the alkali metals, and copper, 
silver, and gold. The alkali metals, so called because their 
hydroxides are strong alkalies, are lithium, sodium, potassium, 
rubidium, and caesium. They are all light metals, with low 
melting points ; while copper, silver, and gold have high 
densities and high fusing points. All of the elements of the 
first group unite with the halogens atom for atom; as, for ex- 
ample, LiCl, Nad, KOI, EbCl, CsCl, Cu 2 Cl 2 , AgOl and AuCl. 
There are also the chlorides 0uCl 2 , AuCl 2 and Au01 3 . 



The Alkali Metals. 





Atomic 
Weight. 


Density. 


Atomic 
Volume. 


Melting 
point. 


Lithium 


7 


0.59 


11.9 


180° 


Sodium 


... 23 


0.97 


23.7 


95.6 


Potassium 


39.1 


0.87 


45.5 


62.5 


Rubidium 


85.5 


1 52 


56.2 


38.5 


Caesium 


.... 133 


1.88 


70.7 


26.5 



The alkali metals constitute a well-defined natural group. 
They all decompose water rapidly at common temperature, 
liberating hydrogen, and quickly tarnish in air. They com- 
bine with the halogens to form compounds containing one 
atom of metal to one atom of halogen, and are therefore 
monads. We have seen that the halogens are strongly acidic, 
and shall shortly learn that the alkali metals are strongly 
basic. They are, in fact, the most basic elements known. 
Potassium, rubidium, and ca3sium have chemically the closest 



84 THE FIRST GROUP. 

resemblance. Lithium approaches in some properties the 
metal magnesium of the second group. Sodium is inter- 
mediate in properties between potassium and lithium. 

The physical properties of the alkali metals vary with their 
atomic weights. Thus their densities and atomic volumes 
increase with their atomic weights, while the melting points 
diminish, as shown in the foregoing table. 

Atomic Volume. — The quotients obtained by dividing the 
atomic weights by the densities (water = 1) of the elements 
in the solid form are called atomic volumes. They represent 
the relative volumes occupied by atomic quantities in the 
solid form. t Thus the atomic volume of lithium is 7 -=-0.59 = 
11.9 ; of caesium it is 133 -=- 1.88 = 70.7. In other words, 7 
grams of lithium measure 11.9 cc, and 133 grams of caesium 
measure 70.7 cc. 



Lithium, Li. 

Atomic Weight, 7. Density, 0.59. 

Lithium was discovered in 1807 by Arfvedson. A number 
of minerals contain a few per cent of it, and a larger number 
contain traces of it. It is widely distributed in the animal 
and vegetable kingdoms, and is found in sea, river, and spring 
waters. It has been detected in meteorites. 

Metallic lithium is obtained by decomposing the molten 
chloride with the current from four to six cells of a Bunsen 
battery. The metal cannot be prepared by heating the car- 
bonate with charcoal. 

Lithium possesses the lustre of silver, is softer than lead, 
but harder than sodium. It tarnishes on exposure to air, but 
not so readily as sodium and potassium. It is the lightest 
known solid. The lithium compounds impart an intense red 
color to the flame, but the metal itself burns in air with a 



SODIUM. 85 

brilliant white light. Lithium attacks glass, porcelain, and 
silica below 200°. 

Lithium Chloride, LiCl, is formed when the metal burns in 
chlorine, and when lithium carbonate is dissolved in hydro- 
chloric acid. The anhydrous chloride separates on evaporat- 
ing a solution of the salt above 15°. 5. At temperatures below 
10° lithium chloride crystallizes with water, forming the 
hydrate LiCl + H 2 0, or LiCl + 2H 2 0. Lithium chloride is 
very soluble in water and alcohol, and is very deliquescent. 



Sodium (Natrium), Na. 

Atomic Weight, 23. Density, 0.97. 

The most abundant compound of sodium- is common salt. 
Sodium is very widely distributed, occurring in all soils, many 
minerals, and in animals and plants. It is so commonly 
present that it is difficult to obtain substances in which the 
spectroscope will not show traces of it. 

Metallic sodium possesses a silver-white lustre, is ductile at 
0°, and soft at ordinary temperature, and melts at 95°. 6. It 
is volatile at a red heat, but attempts to determine its gas 
density have not given satisfactory results, owing to the fact 
that the metal attacked the containing vessels at high temper- 
atures. Sodium was first isolated by Davy in 1807, by the 
electrolysis of sodium hydroxide. 

The metal is now manufactured by reducing sodium car- 
bonate with coal at an intense heat ; sodium distils from the 
retort, and its vapor is condensed in a suitable receiver. The 
reaction is as follows : 

Na a OO a + 2C = 2Na -f 300. 

In a new process, said to give a better yield, an intimate mix- 
ture of finely divided iron ami charcoal are substituted for coal. 



86 THE FIRST GROUP. 

Sodium does not oxidize in perfectly dry air, but tar- 
nishes quickly when a fresh surface is exposed to ordinary air, 
and the metal is slowly but completely corroded. Hence it 
must be preserved from the air either in petroleum oil, by a 
coating of paraffine, or in closed vessels. When the metal is 
placed in hot water the escaping hydrogen is inflamed by the 
heat of the reaction. The decomposition of water by sodium 
at ordinary temperature has already been noticed. The metal 
does not lose its lustre in perfectly dry chlorine gas, but if the 
latter is moist, sodium chloride is found. 

Sodium Chloride, Common Salt, NaCl or Na-Cl, is obtained 
from sea water, salt springs, and salt mines. The salt of com- 
merce is in the form of a fine crystalline meal, such as the fine 
table and dairy salt ; in coarse crystals which form on slow 
evaporation of brine, and in large lumps which come from the 
salt mines. Sodium chloride crystallizes in transparent cubes, 
having a density of 2.16. It is but little more soluble in hot 
than in water of ordinary temperature, as shown by the fol- 
lowing table giving the quantities of sodium chloride which 
are dissolved by 100 parts of water at different temperatures : 

Temperature. 0° 14° 40° 80° 100° 

NaCl 35.5 35.9 36.6 38.2 39.2 

The hydrate NaCl -f- 2H 2 separates on cooling a saturated 
solution of salt to — 10°, and on lowering the temperature to 
- 22° crystals of NaCl + 10H 2 O form. Sodium chloride 
fuses at 776°, and at higher temperatures volatilizes. 

The common salt of commerce contains small quantities of 
sodium and calcium sulphates and magnesium chloride. The 
latter renders salt liable to become damp in the air. 

Pure sodium chloride may be obtained by passing hydro- 
chloric acid gas into a saturated solution of common salt. 
Sodium chloride will separate, and may be washed with con- 



POTASSIUM. 87 

centrated hydrochloric acid and then dried. The impurities 
will remain in solution. 

Exp. 50. — To illustrate the insolubility of sodium chloride in hydro- 
chloric acid add the concentrated acid to a saturated solution of salt. 

The other sodium halides are Sodium Fluoride, NaF ; 
Sodium Bromide, NaBr ; and Sodium Iodide, Nal. 



Potassium (Kalium), K. 

Atomic Weight, 39.1. Density, 0.87. 

Potassium is widely diffused in nature, but is much less 
abundant than sodium. It is found in many minerals and rocks, 
and in all good soils, since its presence is essential to the growth 
of plants. -. Formerly wood ashes were the source of potassium 
compounds. Large quantities of potassium chloride, and an- 
other mineral containing this salt and magnesium chloride, are 
mined near Stassfurth, in Germany. Potassium salts occur in 
small quantity in sea and spring waters. 

Potassium was first isolated by Davy in 1807 by electrolysis of 
potassium hydroxide. Previous to this time the alkalies, sodium 
hydroxide and potassium hydroxide, were regarded as simple 
substances. Lavoisier had, however, supposed that they con- 
tained oxygen from their chemical similarity to well-known 
metallic oxides. The manufacture of potassium by reducing the 
carbonate with coal is more difficult than that of sodium, o wing- 
to the liability of the formation of a black explosive body. 

Potassium is a silver-white metal, brittle at 0°, soft as wax 
at ordinary temperature, and molting at 62°.5. It may be dis- 
tilled in an atmosphere of hydrogen at a faint-red heat. Its 
vapor is green. 



88 THE FIRST GROUP. 

Potassium combines with other elements more energetically 
than sodium. It decomposes water rapidly, and the hydrogen 
evolved is ignited by the intense heat of the reaction, and burns 
with a violet flame. The metal tarnishes immediately in or- 
dinary air, and is soon converted into a mixture of potassium 
hydroxide and carbonate. It, however, does not lose its lustre 
in perfectly pure, dry air. Potassium is a strong reducing 
agent, and has been used to separate other elements from their 
compounds; but sodium being much cheaper, is more generally 
used. Sodium and potassium unite to form an alloy which is 
liquid at ordinary temperature and has the appearance of 
mercury. 

Potassium Chloride, KC1, or K-Cl, crystallizes in cubes, 
has a stronger saline taste and is more soluble in water than 
sodium chloride, which it resembles in properties. It is used 
in the preparation of other potassium salts and in fertilizers. 

Potassium Bromide, KBr, is a valuable medicine. It has a 
sharp saline taste, is readily soluble in water, and crystallizes 
in cubes. 

Potassium Iodide, KI, is a very soluble salt, which forms 
opaque cubes when deposited from hot solutions, and trans- 
parent crystals when a dilute solution is slowly evaporated. 
It is much used as a medicine. Its vapor density has been 
found to be 84.5, theory requiring 83 for the formula KI. 

Potassium Tri-iodide, KI 3 , is obtained by saturating a strong 
solution of potassium iodide with iodine, and evaporating over 
sulphuric acid. It loses two atoms of iodine at 100°. 

Potassium Fluoride, KF, is a very deliquescent salt, obtained 
by neutralizing potassium carbonate with hydrofluoric acid. 
It separates in cubes when the solution is evaporated. 



RUBIDIUM — CAESIUM. 89 

Hydrogen Potassium Fluoride, or Acid Potassium Fluoride, 

HFKF, separates in quadratic crystals, when a solution of 
potassium fluoride in aqueous hydrofluoric acid is evaporated. 
It gives off hydrofluoric acid at a dull-red heat, leaving potas- 
sium fluoride. 



Rubidium, Kb. 

Atomic Weight, 85.5. Density, 1.52. 

Rubidium was discovered by Bunsen in 1861. It is widely dis- 
tributed, but has been found only in minute quantities. It occurs in 
many spring waters, in sea water, and has been detected in plants. 
Metallic rubidium is obtained by the same process as that used for 
sodium and potassium. It is silver-white, and at — 10° is soft. It is 
remarkable for its low melting point, 38°. 5. It oxidizes at once on ex- 
posure to air, and ignites more readily than potassium. It burns like 
potassium on water. The rubidium salts closely resemble the correspond- 
ing potassium salts, and the salts of the two metals crystallize in similar 
forms. 

Rubidium Chloride, RbCl, crystallizes in cubes and is very soluble in 
water. 



Caesium, Cs. 

Atomic Weight, 133. Density, 1.88. 

Caesium was discovered by Bunsen in 1860. It is a very rare element, 
often occurring with rubidium. It is found in some spring waters and 
a small number of minerals, from one of which, lepidolite, it is usually 
prepared. Its salts are not taken from soil by plants. Its compounds 
present great similarity to those of potassium and rubidium. Metallic 
caesium cannot be prepared by heating the carbonate with charcoal, but 
it has been obtained by the electrolysis of a molten mixture of 4 parts of 
barium cyanide and 1 part of coesium cyanide. The metal is soft. 
burns on water like potassium, and quickly ignites on exposure to air. 
Its melting point, 26°. 5, is the lowest of any metal excepting mercury. 

Caesium Chloride, CsCl, is a salt which crystallizes in cubes and de- 
liquesces in damp air. 



SPECTRAL ANALYSIS. 

The vapors of many substances when heated so as to become 
self-luminous appear as bodies of characteristic colors. The 
necessary high temperature may in general be most conven- 
iently secured by the electric spark, but in many cases, nota- 
bly in the alkali and alkali-earth metals, the moderate tern- 
perature of the Bunsen flame is quite sufficient to cause the 
vapors to glow. Thus, sodium, either metallic or in combina- 
tion with other non-metallic elements, will give a remarkably 
pure and brilliant yellow color to such a flame ; while potas- 
sium, so similar in its chemical properties, will glow with a 
characteristic but feeble violet light. This pronounced dif- 
ference in the properties of the two elements has long been 
known, and used as a ready guide to the process of chemical 
analysis of an unknown compound containing one or the 
other of these metals ; only when both sodium and potassium 
are present is there any difficulty in application of the method, 
for then the intensity of the light due to the sodium is so 
great as to quite mask the feeble violet of the potassium, 
unless the latter substance is enormously in excess. In such 
a case chemists were accustomed to observe the flame through 
a plate of cobalt-blue glass, or a film of a solution of indigo, 
both of which are quite opaque to yellow, and very trans- 
parent to the peculiar light of potassium vapor. Thus a 
flame which to the unaided eye appears to be of a strong yel- 
low, indistinguishable from a simple sodium flame, seen 
through a proper thickness of indigo solution might appeal' 
of the characteristic hue of potassium vapor. 

This primitive process of qualitative analysis by the hue of 
a glowing vapor has undergone enormous extension of recent 
years, and yielded many discoveries of the greatest interest. 
Perfected in the hands of Bunsen and Kirchhorf, it has led 



SPECTRAL ANALYSIS. 91 

to the discovery of seven previously unsuspected elementary 
bodies ; to the recognition of the existence of the elements 
familiar to our chemistry throughout the universe, so far as 
the stars are sufficiently bright for examination; to a solution 
of many of the puzzling problems of the phenomena of total 
solar eclipses; to a demonstration of the existence of vast 
cosmical bodies of luminous gases constituting true nebulas; 
and to a method of determining the absolute velocity of motion 
of the fixed stars in the line of vision. 

The fundamental principles of the modern method are as 
follows. If light be passed through a prism, that is, a piece 
of transparent material bounded by two plane polished faces, 
it will be deviated from a direct line towards the thicker side 
of the prism; consequently, since an object appears in the di- 
rection from which its light enters the eye, if one looks 
through such a prism at an object it will appear displaced 
towards the thin edge of the prism. Moreover, since the de- 
viation produced by a prism depends upon the wave-length 
of the light passing through it, and a differing wave-length 
produces a different color sensation, the object will in gen- 
eral appear elongated in the direction of its displacement, 
and variously colored. Thus an ordinary white flame, which 
emits waves of all possible lengths, would appear not only dis- 
placed by a prism, but stretched out into a longer or shorter 
band of color — red on the end nearest the direction of the 
flame, violet at the farther end, and with intermediate colors 
similar in hue and arrangement to those of the rainbow. 

If, instead of a white flame, we observe a Bunsen flame 
tinged by sodium vapor through the prism, we recognize a 
displaced image — wholly unchanged, however, as regards 
color : thus proving that light emitted by sodium vapor con- 
sists of only a single wave-length. Again, if a flame tinged 
by potassium vapor is observed through the prism, wo may 
sec two images, the less deviated being a deep red, ami the 
more deviated a violet: whence wo conclude thai potassium 



92 



SPECTRAL ANALYSIS. 



vapor emits two definite wave-lengths of light — one very long 
and one very short. If sodium and potassium are simultane- 
ously present, we should see three images: a red — least de- 
viated, then a yellow, and last a violet. 

Such images, arranged according to the wave-lengths of 
light emitted by the object, are called spectra, and this 
method of determining the character of the light is called 
spectral analysis. 

If the observed flame has a number of different colors, the 
spectral images belonging to each may overlap, and be recog- 
nized with difficulty unless the flame is a very small one. In 
practice a sufficiently small flame is secured by means of an 




Fig. 65. 



opaque screen with a narrow slit in it, the slit being parallel 
to the refracting edge of the prism. These two elements, 
the prism and the slit, are the only essential elements of the 
apparatus for spectral analysis, and together constitute the 
spectroscope. In practice, however, the spectrum is generallv 
observed through a telescope, thus securing an increase of 
power Avhich otherwise could only be obtained by a greater 
number of prisms. Fig. 65 represents the construction of 
the spectroscope as ordinarily used. Here P is the prism, T 
the telescope, and C a tube bearing the slit, made of adjusta- 
ble width, at its end. The tube C has at its end next the 
prism a lens of focal length equal to its distance from the slit 



SPECTRAL ANALYSIS. 93 

— an addition which admits of the use of a much shorter tube 
for the slit. With this construction the tube C is called a 
collimator. The accessory tube S contains a photographed 
scale at the outer end and a lens at its focal distance from the 
scale at the inner. This is placed at such an angle that when 
the scale is illuminated it is seen by reflection from the face 
of the prism superimposed upon the spectrum, thus aiding in 
the identification of the lines of the various spectra. 

Such a spectroscope, or the simpler prism-slit combination, 
will exhibit a single (or if powerful a closely double) yellow 
line as characteristic of sodium vapor, a deep red and a violet 
line for potassium vapor, a very brilliant red and a feeble 
orange line for lithium ; and so on, no two elements having 
spectra even remotely resembling each other. Since it re- 
quires a very minute quantity of a substance to yield its char- 
acteristic spectrum, a number of rare elements have been dis- 
covered by means of the spectroscope which would otherwise 
have long escaped detection. Such are Caesium, Kubidmm, 
Thallium, Indium, Gallium, Ytterbium, and Scandium. The 
first four of these are named from the color of their brighter 
spectral lines. 

Light from a white-hot solid or liquid body, when analyzed 
by the spectroscope, does not exhibit a system of bright lines, 
but an unbroken band of colors. Such a spectrum is called a 
continuous spectrum. If the light from the white-hot body, 
however, be passed through a metallic vapor and then ob- 
served by the spectroscope, a very remarkable change is rec- 
ognized, provided that the vapor is at a lower temperature 
than the source of light. In this case the continuous spec- 
trum is interrupted by dark lines corresponding exactly in 
place to the blight lines of the metallic vapor, which would be 
visible without the bright source of the continuous spectrum. 
For example, if a sodium- tinted Bunsen flame be observed by 
a spectroscope, a bright-yellow line will be seen ; but it' behind 
the flame a lime-light is placed, a- dark line will appear m the 



94 THE FIRST GROUP. 

yellow of the continuous spectrum. The direct light of the 
sun exhibits thousands of such dark lines when examined by 
the spectroscope, one of them falling exactly in the place of 
the sodium line. Since this dark line is also seen double 
with a spectroscope sufficiently powerful to show the sodium 
line double, the coincidence still being perfect for each com- 
ponent, the conclusion that sodium vapor exists somewhere 
between us and the source of the white light of the sun is in- 
evitable. In the case of iron there are several hundred dark 
lines in the solar spectrum, corresponding, each one, with a 
bright line in the iron spectrum. That the metallic vapors 
are in the sun itself, follows from the observed fact that 
neither the distance of the sun nor the thickness of the layer 
of air through which its light reaches us modifies the essen 
tial characteristics of the spectrum. The same conclusion 
follows from the similarity of the spectra of the light of the 
planets, since all are seen by reflected sunlight, and the gen- 
erally dissimilar spectra of the fixed stars. 

By a comparison of the bright-line spectra of the various 
elements and the solar spectrum, physicists have been able to 
establish the presence of the following elements in the atmos- 
phere of the sun : Hydrogen, Iron, Sodium, Calcium, Mag- 
nesium, Titanium, Nickel. 



Copper (Cuprum)^ Cu. 

Atomic Weight, 63.3. Density, 8.9. 

Copper was known in prehistoric times; occurring native, it 
was available for tools and weapons before the art of iron 
smelting was learned. Native copper is one of the sources of 
the metal in commerce, and sometimes occurs in large masses. 
Copper is extracted from a variety of ores, of which may be 
mentioned the red oxide, Cu 2 0, copper pyrites, CuFeS 2 , and 



COPPEE. 05 

chalcocite, Cu a S. Copper is easily separated from its salts by 
electrolysis, and from its oxides by heating with hydrogen or 
charcoal. The chemistry of the extraction of copper from its 
ores cannot be profitably studied without a knowledge of 
copper compounds. 

Copper is red, tough, moderately hard, capable of a high 
polish, and, including its alloy, brass, ranks next to iron, the 
most useful of metals. It may be rolled into thin leaves and 
drawn into very fine wire. When hammered it becomes stiff 
aud brittle; by heating to redness it is annealed, whether 
cooled quickly or slowly. Copper ranks next to silver as a 
conductor of heat and electricity. Its melting point is 
above that of silver and gold, and it is volatile at very high 
temperatures. Commercial copper usually contains traces of 
silver, arsenic, bismuth, iron, and other metals. It is, how- 
ever, pure enough for most chemical purposes. 

Copper is not affected by moist air free from carbon dioxide, 
nor by dry air, but in the ordinary atmosphere it slowly ac- 
quires a thin green coating of basic carbonate. The metal 
decomposes water slightly at very high temperatures. Strong- 
ly heated in air it oxidizes at once, first to red oxide and then 
to black oxide. In presence of air it slowly dissolves in hydro- 
chloric and sulphuric acids and alkalies. The best solvent for 
the metal is nitric acid. 

Copper forms two well-defined classes of compounds desig- 
nated as cuprous and cupric. One atom of copper replaces 
one of hydrogen to form cuprous compounds; cuprous copper 
is, therefore, apparently a monad. The observed gas density, 
however, of cuprous chloride is 108, corresponding to the 
molecular formula Cu,Cl.,, which requires 98.8. This shows 
that the double atom, Cu,, is a dyad. In no case has the gas 
density of a cupric compound been observed; hence the mo- 
lecular weights of cupric salts are not certainly known. One 
atom of copper replaces two of hydrogen in acids to form 
cupric salts. Copper is therefore regarded as a dyad, as, for 



96 THE FIKST GROUP. 

CI 

example, in cupric chloride, Cu < ™ In cuprous compounds 

the two atoms of copper are supposed to be united by one 
bond, as in cuprous chloride, Cl-Cu-Cu-Cl. 

Cuprous Chloride, Cu 2 Cl 2 or Cl-Cu-Cu-Cl, remains as a 
brown liquid when cupric chloride is heated to redness, and 
on cooling it solidifies to a brown mass. Cuprous chloride is 
usually prepared by boiling a solution of cupric chloride with 
metallic copper. If the solution contains much hydrochloric 
acid cuprous chloride remains dissolved, while in the absence 
of an excess of acid it separates as a white powder. Cuprous 
chloride absorbs moisture and oxygen from the air, and is 
converted into a green oxychloride. Solutions of cuprous 
chloride also take up oxygen and become green; in presence 
of hydrochloric acid the change may be represented as follows: 

Cu Q Cl 2 + + 2HC1 •= 2CuCl 2 + H 2 0. 

Exp. 51.— a. Dissolve a little copper scale in hot hydrochloric acid, 
and pour the solution obtained into water. The formation of a white 
precipitate of cuprous chloride is evidence that the scale contains cu- 
prous oxide. 

b. Place in a test-tube copper scale, copper filings or fine wire, and 
hydrochloric acid, and boil the mixture gently until the green color 
due to cupric chloride disappears. Cuprous chloride will separate if 
but little hydrochloric is present, and the clear hot solution will deposit 
more of the salt on cooling. Pour some of the solution into a porcelain 
dish, and some into water. Note the changes Formulate the reactions 
between the cuprous oxide and hydrochloric acid, the cupric oxide and 
the acid, and also between the cupric chloride and metallic copper. 

Cuprous Bromide, Cu 2 Br 2 , is formed when heated copper is placed in 
bromine. 

Cuprous Iodide, Cu 2 I 2 , is formed by the direct union of copper and 
iodine. Potassium iodide precipitates from a solution of cupric sulphate 
cuprous iodide mixed with iodine : 

2CuS0 4 + 4KI = 2K 2 S0 4 + CuJ a +I 2 . 



SILVER. 97 

CI 
Cupric Chloride, CuCl 2 or Cu< cl , is formed when copper is 

burned in an excess of of dry chlorine. It is a yellowish-brown 
powder, which loses half its chlorine when heated as already 
mentioned. A solution of cupric chloride is best prepared by 
dissolving the oxide or carbonate in hydrochloric acid. On 
evaporation fine green crystals of CuCl 2 + 2H 2 separate. 
The concentrated solution of cupric chloride is green, and 
changes to a bluish-green color on adding much waten 

Exp. 52. — Dissolve cupric oxide, free from cuprous oxide, in hot con- 
centrated hydrochloric acid. Dilute largely with water, and compare 
the results with those obtained with cuprous chloride. 

Cupric Bromide, CuBr 2 , is obtained by dissolving cupric oxide in hydro- 
bromic acid. 

Cupric Iodide, Cul 2 , exists in the dilute solution containing ^ per cent 
of potassium iodide, to which a 1 per cent solution of cupric sulphate 
has been added. Such a solution remains clear, and does not react for 
iodine with starch. 



Silver (Argentum), Ag. 

Atomic Weight, 107.9. Density, 10.5. 

Silver occurs in the free state in nature, and has been known 
from the earliest times. Next to gold it has long been the 
most highly prized of metals, both on account of its beauty 
and valuable properties. It has been found in large masses 
— one in Norway of five hundred and one in Peru of eight 
hundred pounds. Native silver usually contains copper and 
gold, and often mercury, platinum, antimony, and bismuth. 
The larger part of the silver of commerce is obtained from 
ores. Of these the most important are silver sulphide and 
its combinations or mixtures with other metallic sulphides. 
Galena (lead sulphide) always contains silver. Silver chloride, 
bromide, and iodide are also worked as silver ores. Silver 
7 



98 THE FIRST GROUP. 

occurs in sea water to the extent of 1 milligram in 100 liters, 
and has been found in the ashes of plants. 

Polished silver possesses a brilliant white lustre. The surface 
of the metal deposited in the electro-plating process is frosted 
white, while finely divided silver from the reduction of silver 
salts by zinc or copper is gray or even black. Very thin films 
of silver appear blue by transmitted light. Silver ranks in 
malleability next to gold; it may be beaten into leaves yowToo 
of an inch in thickness, and one grain may be drawn into 400 
feet of wire. Silver melts at about 1000°, and may be dis- 
tilled in a lime crucible by means of an oxyhydrogen flame. 
In contact with air molten silver absorbs 22 times its vol- 
ume of oxygen gas, most of which escapes when the metal 
solidifies with a "spitting" or throwing out of particles of the 
silver. The " spitting" is prevented by covering the silver 
with a fusible slag to keep it from contact with the air. Silver 
retains its lustre in pure air; the blackening of silver ware is 
supposed to be due to hydrogen sulphide, which forms a thin 
film of black silver sulphide on the metal. Burnished silver 
retains its whiteness better than a surface which has been 
cleansed by chemicals. Silver is not attacked by dilute sul- 
phuric acid, but dissolves in the hot strong acid. Nitric acid 
somewhat diluted is the best solvent for silver, dissolving it 
readily at ordinary temperatures. 



Exp 53. — Dissolve a quarter of a dollar or other silver coin in the 
smallest possible quantity of nitric acid diluted with its bulk of water. 
The solution contains silver nitrate and cupric nitrate. Dilute with 
about 100 cc. of water, and place in the solution a coil of bright copper 
wire. The silver will precipitate on the wire as a gray spongj^ mass, 
often showing bright crystalline facets. After a time shake it from the 
wire, and when no more is deposited remove the wire, and wash the 
silver by decantation until the washings are colorless. Add some 
ammonia, which will dissolve a little copper adhering to the silver with 
formation of a blue color. Wash again until the washings are colorless. 
Keep the silver for future experiments. 



SILVEH. 99 

Silver Alloys. — The alloys of silver and copper are used for 
coins and solid silver articles, as pure silver is too soft to with- 
stand the wear. The copper imparts toughness and hardness, 
and unless present in too large proportions the alloy has the 
whiteness and beauty of pure silver. The United States silver 
coinage consists of 900 parts of silver to 100 of copper, and 
the British (sterling silver) contains 925 parts of silver to 75 
of copper. 

Silver alloys readily with many other metals, but the com- 
pounds formed are of little importance, except those of lead 
and zinc, which are obtained in some processes of separating 
silver from its ores. 

Silvering and Silver Plating. — A coating of silver on other 
metals is highly valued because of the beauty of the silver sur- 
face and its resistance to corrosion. Copper, brass, and Ger- 
man silver articles may be silvered by first rubbing on a silver 
amalgam, heating to volatilize the mercury, and then burnish- 
ing. Or they may be rubbed with or put into a solution of 
a silver salt, when a thin film of silver will be deposited on 
them. These methods do not answer for articles which are 
subjected to much wear. In order to obtain a more durable 
coating they are electro-plated with silver. In this process the 
piece to be plated is placed in a solution of silver cyanide in 
potassium cyanide, and connected with the negative pole of a 
battery, and the positive pole is joined to a sheet of pure 
silver suspended in the solution. The galvanic current de- 
composes the silver cyanide, and the silver is deposited on the 
metallic piece as a firm, compact coating. At the same time 
silver passes into solution from the sheet of silver, thus 
maintaining the strength of the solution. 

Silver may be deposited as a brilliant mirror on glass from 
alkaline solution to which milk-sugar or certain other organic 
substances have been added. 

The molecular weights of silver salts are not certainly known. 



100 THE FIRST GROUP. 

as the gas density of no one of them has been observed. One 
atom of silver replaces one atom of hydrogen in acids to form 
the common or argentic salts, and silver is therefore regarded 
as a monad. The formulas of the argentic salts are conse- 
quently analogous to those of the alkali metals. The argentous 
compounds are little known. 

Silver Fluoride, AgF. — Silver oxide dissolves in hydrofluoric 
acid, and the solution yields on evaporation deliquescent crys- 
tals having the composition AgF + 2H 2 0. 

Silver Chloride, Argentic Chloride, AgCl, or Ag-Cl, is pre- 
pared by adding common salt or other soluble chloride to solu- 
tions of silver nitrate or sulphate. The white curdy precipi- 
tate on long standing or violent agitation settles as a heavy 
powder, leaving the liquid clear. When silver chloride is 
exposed to light it first turns violet and finally black, with loss 
of a little chlorine. It is extremely insoluble in water, and 
but very slightly soluble in concentrated nitric acid. One 
part of silver chloride dissolves in 200 parts of concentrated 
hydrochloric acid. It is somewhat soluble in solutions of 
alkali chlorides, and is very easily dissolved by ammonia water, 
by alkali thiosulphates and c}^anides. Silver chloride melts at 
260° to a yellow liquid, which cools to a brown, tough, and 
sectile mass. If moist silver chloride is placed in contact with 
iron or zinc the silver gradually separates in a spongy form, 
and the chlorine unites with the other metal. 

Silver Bromide, AgBr, is very similar in properties to silver 
chloride, from which it differs by being less soluble in ammonia, 
and having a yellow tinge as it is ordinarily precipitated. It 
is used in various photographic processes. 

Silver Iodide, Agl. — Silver dissolves in concentrated hydriodic 
acid, setting hydrogen free. From the hot saturated solution 



GOLD. 101 

crystals form on cooling, which contain silver iodide and hy- 
driodic acid. Silver iodide separates from a solution of silver 
nitrate on the addition of soluble iodides as a bright-yellow 
curdy precipitate. Pure silver iodide is changed very slowly, 
or not at all, by sunlight, but if mixed with silver nitrate it 
acquires a greenish-black color. Silver iodide resembles the 
chloride and bromide, but is very slightly soluble in ammonia. 

Exp. 54. — Convert about half of the silver of the previous experiment 
into silver nitrate, AgN0 3 , by dissolving in nitric acid. Dilute a por- 
tion of the solution largely with water, add hydrochloric acid as long as 
a white precipitate forms, and agitate violently until the chloride settles 
quickly. Wash the precipitate on a filter until free from acid. Expose 
a portion of the silver chloride to sunlight. Treat another part with 
strong ammonia water. The remainder of the silver chloride may be 
rinsed into a test-tube, and a piece of zinc placed in contact with it. 
Note the results. 

Exp. 55. — Add some of the silver nitrate of the foregoing experi- 
ment to a solution of potassium bromide in a test-tube. Shake vio- 
lently, and after the precipitate has subsided decant the supernatant 
liquid, and then add ammonia water to the precipitate. Repeat the ex- 
periment, using potassium iodide instead of potassium bromide. For- 
mulate the reactions between the silver nitrate and the potassium bro- 
mide and iodide. 



Gold (Aurum), Au. 

Atomic Weight, 19G.7. Density, 19.3. 

Gold was prized in prehistoric times for its beauty, its dura- 
bility, and ease of working. It usually occurs in the free 
state, sometimes crystallized in octahedral and dodecahedral 
forms, but usually in irregular nuggets, flattened scales, or 
"fine dust." It is widely distributed, though found in com- 
paratively few localities in paying quantities. Most river 
sands contain gold. Traces of gold are found in copper and 
iron pyrites, galena, and in many silver ores, ami a very 



102 THE FIRST GEOUP. 

minute proportion occurs in sea water. Native gold always 
contains silver, and sometimes minute quantities of copper, 
iron, palladium, and rhodium. 

Pure gold in mass has a bright-yellow color, is somewhat 
harder than lead, and is the most malleable and ductile of 
metals. Gold-leaf may be beaten out so thin that 280,000 
leaves together have but an inch of thickness, and one grain 
of gold has been drawn into a wire 500 feet long. The thin- 
nest gold-leaf and a thin deposit of gold upon glass appear 
green by transmitted light. Finely divided gold suspended 
in a liquid has a red color by reflected and a blue color by 
transmitted light. 

Gold melts at about 1240°, and at the highest temperature 
of a wind-furnace it volatilizes slightly. Gold does not oxi- 
dize or tarnish in the air at any temperature, and is not 
altered by sulphur, which blackens silver. It is not acted 
upon by any single acid, excepting selenic, under ordinary 
circumstances. When, however, gold immersed in sulphuric 
acid is made the positive electrode, of a galvanic battery, it 
dissolves, and in nitric acid under like conditions a violet 
precipitate forms. Finely divided gold dissolves when heated 
with concentrated sulphuric acid, mixed with either iodic 
acid, potassium permanganate, or manganese dioxide. 

The ordinary solvent for gold is a mixture of hydrochloric 
and nitric acids. Solutions containing free chlorine or bro- 
mine attack it ; iodine has slight action. 

Gold Alloys. — Gold mixes in fusion with most metals, but its 
important alloys are those with silver and copper. These alloys 
are harder and more useful than pure gold. The proportion 
of gold in jewelry is expressed in carats, 24 carats being pure 
gold, and the fineness of bullion is represented in parts per 
thousand. United States gold coin is 900 fine or 21.6 carats, 
and British coin is 916. GQ or 22 carats, fine. The alloys of gold 
and silver have a lighter color than pure gold, and with certain 



GOLD. 103 

proportions of silver are of a greenish-yellow color. Copper 
imparts a reddish tinge. The alloys used for jewelry contain 
both copper and silver, in varying proportions according to the 
color desired ; 14-carat gold possesses good color, and is well 
adapted for watch-cases and trinkets. 

Refining. — "Fine gold" is obtained from bullion and alloys 
by " parting," i.e., dissolving out the base metal with hot 
nitric acid or sulphuric acid. If more than 25 per cent of 
gold be present, the alloy is melted with sufficient silver to 
reduce the gold to that proportion, as richer alloys do not 
leave pure gold when treated with acids. 

Exp. 56. — a. Alloy, by melting together on charcoal, a gold dollar 
with 3 parts pure silver, or silver coin. Hammer or roll out the 
button into a thin strip, annealing from time to time by heating to dull 
redness. Roll the strip into a compact coil, and boil it with nitric acid 
of density 1.30. After the action has ceased decant off the acid solu- 
tion, and boil with a fresh portion of acid. Wash by decantation, adding 
the first washings to the nitric acid solutions, from which the silver 
may be recovered by precipitation as chloride. See Exp. 54. 

b. Take about one fourth of the gold, dry, and fuse it on charcoal 
with the blowpipe. Hammer out the globule of fine gold somewhat, 
and anneal. 

c. Dissolve the remainder of the gold in a mixture of 3 parts of con- 
centrated hydrochloric and 1 part of nitric acid. The solution contains 
hydrogen auric chloride, HAuCl 4 . Evaporate the solution nearly to 
dryness, add 100 cc. of water, and reserve for further experiment. 

Gold is also refined by the following method : 
The alloy — coin for example — is f nsed under borax, and a 
stream of chlorine gas is passed into the molten metal. The 
metals, other than gold, are converted into chlorides, which 
may be poured off as soon as the gold cools and solidities. 

None of these refining processes yield entirely pure metal, 
and, in order to remove the last traces of silver, the gold is 
dissolved in nitn) hydrochloric acid, and precipitated from the 
solution by oxalic acid or ferrous sulphate. 



104 THE FIRST GROUP. 

Exp. 57.— a. Take 10 cc. of gold solution, add a small crystal of 
oxalic acid, then potassium carbonate till the solution becomes clear, 
then an excess of oxalic acid, and boil ; the gold quickly separates in 
yellow spangles. 

b. Dissolve a bit of ferrous sulphate, half the size of a pea, in half a 
test-tube of cold water. Add two or three drops of gold solution to half 
a test-tube of water, and mix the two liquids. The gold will be precipi- 
tated in so fine a state as to remain suspended in the liquid, and will 
appear red by transmitted light and bluish by reflected light. 

Amalgam. — Gold unites readily with mercury, and appar- 
ently dissolves in an excess of the latter metal, but on straining 
the mixture through wet leather a solid amalgam is obtained, 
and the strained mercury contains only a trace of gold. A 
very small quantity of mercury is sufficient to whiten the 
surface of gold. 

Gilding. — Gold-leaf is used for gilding wood, paper, and 
leather, being made to adhere by means of glue or varnish. 
Metals were formerly gilded by rubbing with a gold amalgam, 
the mercury being then expelled by heat, or they were dipped 
into a solution of gold ; but these methods are largely replaced 
by the electro-plating process. The amount of gold required 
to cover a metallic surface is exceedingly small. China and 
glass are gilded by applying precipitated gold (Ftp. 57, b), with 
oil of turpentine, fusing upon the surface of the ware, and 
burnishing. 

Gold Purple. — A solution of stannous chloride containiug 
a little stannic chloride produces in solutions of gold chloride 
a purple-red precipitate which sometimes inclines to violet or 
brownish red, and has long been known as Purple of Cassius. 
This precipitate is insoluble in hydrochloric acid. Its com- 
position is doubtful, and in the moist state it does not appear 
to contain metallic gold, as it is soluble in ammonia. It does 
not give off oxygen on heating, a fact due perhaps to the 



GOLD. 105 

stannous compounds in it taking up any oxygen separated by 
heat from the gold. 

A gold purple is also obtained by adding an excess of mag- 
nesia to a solution of gold chloride. In this case the gold is 
precipitated as hydroxide on the magnesia. The magnesium 
chloride formed is removed by washing with water, and the 
mixture of gold hydroxide and magnesia is ignited. The 
color of the product varies with the proportion of gold pres- 
ent ; 0. 1 per cent producing a faint tint, and 5 per cent a fine 
purple. In place of magnesia the gold may be precipitated 
on other substances, and a similar purple obtained. The color 
appears to be- due to finely divided gold which reflects red 
light. Gold purple is used to give a fine red color to glass 
and porcelain. 

Exp. 58. — Boil a piece of tin foil with hydrochloric acid until it is 
mostly dissolved, pour off the solution of stannous chloride, and add to 
it a drop or two of chlorine water to convert some of the stannous 
chloride into stannic chloride. Finally, mix with a dilute solution of 
gold. 

Aurous Chloride, AuCl, is prepared by heating chlorauric 
acid or auric chloride to 185°. It is a yellowish powder, which, 
in contact with water, decomposes into gold and auric chlo- 
ride : 

3AuCl = 2Au + AuCl 3 . 

Gold Dichloride, Auroso-Auric Chloride, AuCl 2 or Au 2 Cl 4 , is 
prepared by acting on spongy gold with dry chlorine gas. It 
is a dark-red hygroscopic substance. It decomposes in pres- 
ence of water into aurous and auric chlorides, which may be 
separated by quick filtration. The aurous chloride by pro- 
longed action of water is further resolved into gold and the 
trichloride. 

Auric Chloride, AuCl 3 . — Gold trichloride is formed when 
chlorine is passed over gold-leal' heated to 820-230°. It may 



106 THE FIRST GROUP. 

be sublimed in a stream of chlorine and obtained in reddish 
crystals. A solution of AuCl 3 is prepared by acting on gold 
dichloride with water, when metallic gold is separated and 
Au01 3 is formed. The solution of auric chloride is brownish 
red, and yields on evaporation orange-red crystals of Au01 3 -f- 
2H 2 0. 

Chlorauric Acid, Hydrogen Auric Chloride, HAuCl 4 or 
HCl.AuCl 3 . — This substance is known only in combination 
with water. On adding hydrochloric acid to a solution of 
auric chloride the color changes from red to yellow, and the 
solution yields on evaporation yellow crystals of HAu01 4 -f- 
4H 2 The same compound is commonly made by dissolving 
gold in aqua regia (Exp. 56, 6*). The crystals are deli- 
quescent, and on moderate heating lose hydrochloric acid, 
chlorine, and water, and are converted into aurous chloride, 
which on heating more strongly is resolved into metallic gold 
and chlorine. 

Chlorauric acid forms a number of salts, as for instance 
sodium chloraurate, NaAuCl 4 + 2H 2 0. This salt is obtained 
in crystals by adding NaCl to a solution of HAuCl 4 and 
evaporating the solution. 



THE SIXTH GROUP. 

The members of this group are oxygen, sulphur, selenium, 
tellurium, chromium, molybdenum, and uranium. 



Oxygen, 0. 

Atomic Weight, 16. Molecule, 2 . 

Oxygen is a colorless, tasteless, and odorless gas. It occurs 
in the free state in the atmosphere, which is a mixture of 
about one fifth oxygen and four fifths nitrogen by volume. 
It is distinguished from all other gases, nitrous oxide excepted, 
by causing a glowing splinter of wood to inflame. It is the 
supporter of ordinary combustion, and substances which burn 
in air burn with increased brilliancy in oxygen. 

Oxygen was first isolated from other gases by Priestley in 
the year 1774, who obtained it from red oxide of mercury. 
He discovered that when this substance is heated a gas is 
given off which is insoluble in water, and supports combustion 
better than air. Previous to 1774 the air was regarded as a 
mixture of two gases, one of which supports combustion. 

Oxygen is the most widely distributed and abundant ele- 
ment of the earth's crust. It constitutes nearly one half of 
the weight of most rocks and soils, eight ninths of water, and 
a large proportion of the weight of animals and plants. 

Oxygen gas has a density of 1G (hydrogen as unity), ami is 
about one tenth heavier than atmospheric air. It liquefies 
at — 118° at a pressure of 50 atmospheres, and boils at — 181° 
under a pressure of one atmosphere. The density of liquid 



108 



THE SIXTH GKOUP. 



oxygen is 1.124. Oxygen is very slightly soluble in water, 
100 volumes absorbing at ordinary temperature about 3 vol- 
umes of the gas. 

Oxygen is obtained by heating various compounds of it and 
by the electrolysis of water. Potassium chlorate is commonly 
used in the laboratory in preparing oxygen. The salt, when 
heated, is decomposed into potassium chloride and oxygen : 

KOIO, = + KC1 + 30. 

The reaction occurs at lower temperatures, and the evolution 
of gas is more regular if the chlorate is mixed with half its 




Fig 



weight of manganese dioxide. The action of the dioxide is 
not understood. It does not give off oxygen at the tempera- 
ture required to decompose the chlorate of the mixture, and 
remains unchanged mixed with potassium chloride after the 
heating. 



Exp. 59.— Heat a little red mercury oxide, HgO, in a narrow hard- 
glass tube closed at one end. Place a glowing splinter in the escaping 



OXYGEN". 



109 



gas. Kepresent by an equation the decomposition of the mercury 
oxide. 

Exp. 60. — Mix 10 grams of potassium chlorate in crystals with half, 
or better an equal weight of manganese dioxide, and place the mixture 
in an ignition-tube about the size of, but a little heavier than, a common 
test-tube. Connect the ignition-tube by means of a cork with delivery- 
tube having a joint of rubber tubing. Fasten the ignition-tube with 
copper wire to a lamp-stand ring, as shown in Fig. 66, so that the end 
with the cork shall be a little lower than the sealed end. Heat cautiously 
with a lamp held in the hand, and collect the gas over water. When 
4he evolution of gas ceases, remove the delivery-tube from the water. 
Jars containing oxygen may be taken from the water-pan by closing 
them with wet cardboard, and left covered with the cardboard until 
required for use. But little gas will escape in half an hour. Oxygen 
thus prepared will show a slight cloud, and will contain a little chlorine. 
The former will disappear, and the latter will be absorbed by water 
after a time. 

Calculate the weight and number of liters of oxygen which 10 grams 
of potassium chlorate will yield. The weight of one liter of oxygen at 
0° and 760 mm. may be found by multiplying the weight of 1 liter of 
hydrogen by the density of oxygen. 

Exp. 61.— Thrust a glowing splinter of wood into a jar of oxy- 
gen. 

Exp. 62.— Place a piece of charcoal on a chalk spoon, which is sup- 
ported by a copper wire passing through a cardboard 
cover. Ignite a corner of the charcoal and place it in a 
jar of oxygen, as shown in Fig. 67. 

Exp 63. — Anneal a thin watch-spring by heating to 
redness in the lamp flame, and bend it into a spiral. 
Fasten one end into a cardboard cover and stick the 
other end into a piece of match-stick about i of an 
inch long. Ignite the bit of wood, and when it is 
partly burned place the watch-spring in a piut jar of 
oxygen. The burning wood serves to heat the iron to 
the temperature at which it burns in oxygen. 

Exp. 64.— Place a bit of sulphur on a chalk spoon in oxygen. Note 
whether any changes occur. Ignite the sulphur and put again into 
oxygen. Notice odor, and test the gas in the jar after the combustion 
with blue litmus paper. 

Exp. 65. — Bum a small bit of phosphorus in oxygen, and test the 
water in the jar with blue litmus paper. This experiment should not 




Fig. 67. 



110 THE SIXTH GKOUP. 

be attempted by beginners on account of liability to severe burns in 
handling phosphorus. 

Exp. 66. — Lift from water a half-gallon jar of hydrogen, and, hold- 
ing the jar mouth downwards, ignite the gas; then pass a tube from 
which a slow stream of oxygen is issuing into the hydrogen. The tube 
should be tipped with a smaller tube of platinum foil. 

Exp. 67.— a. Heat a piece of platinum wire in a hydrogen flame from 
a compound blowpipe, then turn on oxygen and observe that the flame 
becomes smaller and that the platinum melts. 

The quantity of heat produced by burning a given amount of hydro- 
gen is the same whether the gas burns in air or pure oxygen. When, 
however, it burns in air the flame is largely diluted with and cooled 
by the nitrogen of the air, the nitrogen taking no part in the combustion. 

b. Heat a piece of chalk in the compound blowpipe flame. The light 
thus obtained is known as the calcium light, and is used when a very 
intense light is required. In place of hydrogen, illuminating gas is 
commonly employed for the calcium light. 

c. Melt the end of a file and volatilize a bit of zinc with the compound 
blowpipe flame. The zinc is supported on a piece of charcoal. 

The burning of fuels, illuminating oil and gas is a rapid 
chemical union of the carbon and hydrogen in them with the 
oxygen of the air, carbon dioxide gas and water being formed. 
Oxygen is commonly called the supporter of combustion, but 
fuel is as truly a supporter of combustion, since it is essential 
to a fire, and in this sense oxygen may be regarded as a fuel. 
In Exp. 66 we have seen that hydrogen is apparently the sup- 
porter of the combustion of a jet of oxygen. 

The combining of oxygen with substances is termed oxida- 
tion, and the compounds thus formed are called oxides. Many 
bodies slowly unite with oxygen, especially in moist air. Iron, 
for example, rusts, taking both oxygen and water from the 
air. Oxidation goes on gradually in the decay of animal and 
vegetable substances. All of the elements, excepting fluorine, 
are known to unite either directly or indirectly with oxygen. 

Ozone : Molecule, 3 . — When an electric discharge passes 
through oxygen a pecnliar odor is noticed, and the gas pos- 



OZONE. Ill 

sesses more active properties,, a portion of the oxygen having 
changed to an allotropic* modification known as ozone. 

Under favorable conditions but a few per cent of a given 
volume of oxygen can be ozonized, but by passing oxygen 
containing ozone through a tube cooled to — 181° by means 
of liquid oxygen, pure ozone may be obtained as a steel-blue 
liquid, boiling at — 106° and changing to a blue gas. The 
gas density of ozone has been found to be 24; its molecular 
weight is therefore 48. Ozone gradually reverts to oxygen at 
ordinary temperature, and quickly at 200°. 

Ozone is produced in minute quantities in many cases of 
combustion, and in the electrolysis of water. It differs from 
ordinary oxygen in that it is more active chemically. Silver 
and mercury, which are not affected by ordinary oxygen, are 
at once oxidized by ozone not perfectly free from water, f 
Ozone liberates iodine from potassium iodide, and rapidly 
oxidizes many animal and plant substances. Air containing 
rather more than sufficient ozone to be perceptible by the 
smell, has an irritating effect on the throat. The atmosphere 
contains ozone, but the quantity is so small that it is detected 
with difficulty. 

Exp. 68. — The tube AB, Fig. 68, contains a platinum wire, sealed into 
the end A. About the tube is a coil of copper or platinum wire. The 
wash-bottle C contains concentrated sulphuric acid to dry the oxygen 
which is supplied from a gas-holder. Pass a slow current of oxygen 
through the apparatus and into the jar D, in which is hung a strip of 
paper smeared with starch-paste containing potassium iodide. The 
test-paper will not change color. Connect the inner and outer wires of 
the tube with an induction coil giving a half-inch spark. A silent dis- 

* Allotropy is the property of existing in two or more modifications, 
differing in physical or chemical properties, or both. Thus there are 
the ordinary and ozone forms of oxygen, and the three forms or modi- 
fications of carbon, namely, charcoal, graphite, and diamond. 

f Shenstone and Cundall have found that ozone, dried by long <. \ 
posure to phosphorous pentoxide, reverts to oxygen after contact for 
some hours with mercury, without apparently oxidizing the metal. 



112 



THE SIXTH GEOUP. 



charge of electricity will take place between the inner surface of the 
tube and the platinum wire, and the odor of ozone will soon be per- 
ceptible. The test paper will be colored blue by the liberated iodioe. 
The reaction between the ozone, potassium iodide and water is as fol- 
lows : 

3 + 2KI + H 2 - 2KOH + I 2 + 2 . 

Other substances, chlorine for example, liberate iodine from potassium 
iodide, but in the change potassium salts are formed which are not 
alkaline towards litmus paper. Hence, to show the presence of potas- 
sium hydroxide, KOH, as indicated by the equation, expose to ozone a 
delicate red litmus paper, and one which is saturated with a dilute 
solution of potassium iodide. The former will not change color, 
showing the absence of an alkaline gas ; while the paper with the iodide 




will be colored blue by the potassium hydroxide formed. A similar 
test-paper exposed to chlorine will not turn blue. 

Exp. 69. — Agitate some clean dry mercuiy in a jar of dry oxygen. 
It will not change iu appearance. Next shake up some mercury in a 
jar of ozonized oxygen. The metal will lose its clean, bright surface, 
and will stick to the glass. After a few minutes the gas in the jar will 
not react for ozone with potassium iodide starch-paper. 

Exp 70. — Connect the end of the ozonizing tube (Fig. 6S) bj r means 
of a black rubber tube with a glass tube, and attempt to collect the gas 
over water. In a few minutes the rubber will be perforated by the 
ozone. 

Exp. 71.— Place some pieces of phosphorus in a large jar of air, and 
partly cover them with water to prevent the phosphorus from inflaming. 
After a time test the air in the jar for ozone. 



WATER. 



113 



Exp. 72. — Hang a wet potassium iodide starch paper in a pint jar 
containing a few drops of ether, and introduce a hot glass rod. The 
ether will be oxidized to a pungent body known as aldehyde, and at 
the same time a portion of the oxygen of the air in the jar will be ozo 
nized. 

Hydrogen and oxygen unite to form the compounds hydro- 
gen monoxide or water, H 2 0, and hydrogen dioxide, H 2 2 . 

Water, H 2 or H-O-H. — The abundance of water is familiar 
to all. It forms a large proportion of the weight of animals 
and plants, and is a constituent of many minerals and salts, 
and is always present in the atmosphere. 

Water is a tasteless, odorless liquid, boiling at 100° under 
a pressure of 760 mm. of mercury. It may be cooled below 
0° without freezing, but the temperature of melting ice is 
practically constant, and is not changed to a measurable ex- 
tent by the ordinary changes in atmospheric pressure. Water 
and ice give off vapor when exposed to the air. The following 
table gives the pressure of the vapor of water at different 
temperatures : 



TABLE OF THE PRESSURE OF THE VAPOR OF 


WATER. 


Tem- 


Pressure 


Tem- 


Pressure 


Tem- 


Pressure 


Tem- 


Pressure in 
Atmospheres. 


pera- 
ture. 


in milli- 
meters of 


pera- 
ture. 


in milli- 
meters of 


pera- 
ture. 


in milli- 
meters of 


pera- 
ture. 


1 atmosphere 
= 760 nun. 




mercury. 




mercury. 




mercury. 




of mercury. 


-20° 


0.927 


12 


10.457 


25 


23.550 


100 


1 





4.600 


13 


11.162 


26 


24.988 


111.7 


1.5 


1 


4.940 


14 


11.908 


27 


26.505 


120.6 


o 


2 


5.302 


15 


12.699 


28 


28.101 


127.8 


2.5 


3 


5.687 


16 


13.536 


29 


29.782 


133.9 


3 


4 


6.097 


17 


14.421 


30 


31.548 


144.0 


4 


5 


6.534 


18 


15.357 


40 


54.906 


159.2 


6 


6 


6.998 


19 


16.346 


50 


91.982 


170.8 


8 


7 


7.492 


20 


17.391 


60 


148.791 


180.3 


10 


8 


8.017 


21 


18.495 


70 


233.093 


224.7 


25 


9 


8.574 
9.165 
9.792 


22 
23 
24 


19.659 

20.888 
22.184 


80 

90 

100 


354.280 
525.450 
760. 






10 






11 













114 



THE SIXTH GROUP. 



Exp. 73. — The boiling of water at less than atmospheric pressure 

may be shown by an apparatus 
such as represented by Fig. 69. 
The vessel A has a capacity of 
about a liter. Its neck is sur- 
rounded by the tube D, through 
which water may be passed. The 
thermometer should not touch 
the sides of the neck of A. The 
tube B has a diameter of 8 to 10 
mm., and a height above C of 
800 mm. 

Boil water vigorously in A 
until the air is expelled, allow- 
ing the steam to escape freely 
from the lower end of B, and 
then place a vessel containing 
mercury under B as indicated 
by the figure, and remove the 
lamp used in heating the water. 
To avoid danger in case A should 
crack, it is surrounded by a glass 
jar with card-board cover, as 
shown. The water will continue 
to boil, and the mercury will rise 
in B. On passing a stream of 
water into the jacket-tube D, the 
steam will be rapidly condensed, 
and the water will boil inter- 
mitt ingly until the thermometer 
has fallen to about 30°, when 
Fi g. 69. the column of mercury in B 

will be found to be about 30 mm. shorter than the barometric 

column. 




Pure water has a bluish-green color when viewed in large 
mass. It has the maximum density at 4°, and it expands 
slightly when the temperature sinks below or rises above this 
point. Water in freezing increases in volume nearly ^y, hence 
ice always floats on water. 



WATEE. 115 

Water is the most common solvent, dissolving very many 
salts and other bodies and gases. Solution is a physical 
phenomenon, accompanied in some cases by chemical changes. 
Common salt dissolves unchanged, and may be recovered by 
evaporating the water. Sodium also dissolves in water, part 
of which is decomposed, and the solution yields on evapora- 
tion not sodium, but sodium hydroxide. 

Rain water is nearly pure, but river and spring waters con- 
tain saline matter dissolved out of the soil and rocks through 
which the water has passed. Vegetable matter and products 
of its decay are often serious impurities in drinking water. 
Some spring waters owe their medicinal properties to the 
salts they contain, or to the gases hydrogen sulphide and 
carbon dioxide held in solution. Sea water contains about 
3.5 per cent of solid matter — mostly common salt. For ordi- 
nary laboratory uses water is purified by distilling and con- 
densing the steam in a block-tin pipe. Glass does not answer 
well, as it is slightly soluble in water. 

Exp 74 — Boil tap water in a glass retort, and condense the steam in 
a flask placed over the neck of the retort. The flask may be placed in 
a pan containing water, and cooled by turning occasionally, or covering 
with a wet towel. Notice when heating the water in the retort that air 
bubbles escape before the water boils, and that the steam which fills 
the retort when the water is boiling is invisible. After about two thirds 
' has distilled, evaporate a few drops of the distilled water on platinum 
foil. A faint residue will be perceptible on heating the foil to redness. 
Evaporate as before some of the water remaining in the retort. 

Composition of Water. — Numerous experiments have been 
made with the greatest possible care in order to find the pro- 
portion of oxygen which is contained in water. The weight 
of evidence thus far obtained indicates that 15.96 weights of 
oxygen combine with 2 weights of hydrogen, and that the 
atomic weight of oxygen is a little less than 16, the number 
adopted in this book. 



116 



THE SIXTH GROUP. 



Two volumes of hydrogen combine with one volume of 
oxygen to form two volumes of steam : 



1 




1 


+ 


16 


= 


9 


9 



Hydrogen. 



Oxjgen. 



Steam. 



The gas density of steam is 9; twice the density equals 18, 
the molecular weight of water. In 18 weights of water are 
2 weights of hydrogen, the weight of 2 atoms of hydrogen; 
and 16 weights of oxygen, the weight of 1 atom of oxygen. 
The water molecule is, therefore, composed of 2 atoms of 
hydrogen and 1 atom of oxygen, and has the formula H 2 Oo 
The reaction occurring when the gases unite has commonly 
been represented thus : 



2ET 



B = 2HO. 



Later it will be shown that the change is probably more com- 
plicated. 

The burning of hydrogen with pure oxygen produces the 
intense heat of the compound blowpipe, which Bunsen has 
calculated to be 2844°. The quantity of heat evolved is very 
large, 2 grams of hydrogen burned with oxygen producing 
68,360 calories. 



Exp. 75. — Blow a small soap-bubble with a mixture of 2 volumes of 
hydrogen and 1 volume of oxygen contained in a rubber gas-bag. Re- 
move the bag and apply a flame to the bubble. The shock is due to the 
expansion caused by the intense heat of the combustion, and also to the 
sudden condensation of the steam formed, producing a partial vacuum 
into which the air rushes. There is danger of in juring the ears if the 
large quantities of the mixed gases are exploded. 

Exp. 76.— Remove the flask F, Fig. 70, and ignite a jet of oxygen and 
hydrogen at the blowpipe-tip A, then turn the flame low and replace 



WATER. 



117 




the flask. As soon as the moisture is driven from the upper part 
of it the flame may be made larger. The steam 
formed will he condensed in the cooling-tube 
G When a little water has been collected in the 
beaker, the latter may be held up against the deliv- 
ery-tube, so that escaping gas shall bubble through 
the water; then the supply of oxygen may be regu- 
lated so that little or no gas escapes. The height 
of the water in the gas-holders supplying oxygen 
and hydrogen should now be marked, and again at 
the close of the experiment, In case the flame is 
accidentally extinguished, the gas in the flask F 
should be displaced by a current of air, to make sure 
before relighting that F does not contain an ex- 
plosive mixture of oxygen and hydrogen. 

Exp. 77. — Pass a current of electricity from 4 
or 6 Bunsen cells through water containing ^ by 
volume of sulphuric acid. The apparatus repre- 
sented by Fig. 71 is convenient for this purpose. 
The hydrogen obtained may be tested by burning it, Fig. 70. 

and the oxygen by holding a glowing splinter in the escaping gas when 
the stopcock is opened. If the two gases are measured very accu- 
rately it will be found that for every two volumes 
of hydrogen collected there is slightly less than 1 
volume of oxygen, as the latter is more soluble in 
water than the former, and a very small part of the 
oxygen exists as ozone, as may be proved with a 
potassium iodide starch test-paper. Pure water is 
a poor conductor of electricity, and is not decom- 
posed by it. It is supposed that the sulphuric acid 
in the solution elcctrolyzed is first decomposed into 
hydrogen, oxygen, and S0 3 , thus: 

H 2 SO, =H a + + SO s . 

S0 3 is known to form sulphuric acid with water : 
80 3 + H a O = H 2 S0 4 . Whatever may be the 

changes, the quantity of sulphuric acid in the 
IG " (1 ' solution is not diminished by the electrolysis. 

The best method of obtaininc; a mixture of 2 volumes of hydrogen and 




118 



THE SIXTH GKOTJP. 



1 volume of oxygen is to decompose dilute sulphuric acid by electricit}" 
in the apparatus shown in Fig. 72. The air in the apparatus is expelled 
by the escaping gases, and the quantity of ozone formed is too small to 
interfere with most uses of the mixed gases. 

Exp. 78. — The apparatus represented by Fig. 73 is designed to show 
the volume of the steam which results from the union of known vol- 
umes of hydrogen and oxygen. The tube A is of stout glass, 80 or 90 
cm. in length, and having a diameter of about 12 mm. Two platinum 
wires are sealed into the upper and closed end. These ends within the 
tube are near together, but not in contact. The tube A is heated by 





Fig 



Fig. 73. 



steam passed into the jacket-tube B. The glass cylinder C is nearly 
filled with mercury 

Dry A thoroughly by warming and sucking air through it by means 
of a long glass tube, and while still warm fill it with warm clean mer- 
cury. Adhering bubbles of air may be removed by allowing a large bub- 
ble to pass the length of the tube several times. When A is filled with 
mercury close it with the thumb, and place the open end under the mer- 
cury in the cylinder. Pass steam into the jacket-tube, and allow a mix- 
ture of two volumes of hydrogen and one of oxygen from the apparatus 
shown in Fig. 72 to bubble into A. Mark the volume of the gas by 



HYDKOGEN DIOXIDE. 119 

means of a strip of paper pasted on B, and also mark the height of the 
mercury in A above the mercury in 0. A strip of wood cut the length of 
the mercury column is a convenient measure. Next lower A together 
with B until the end of A rests on a rubber pad in the bottom of the cyl- 
inder, and pass a spark from an induction coil between the wires sealed in 
A. The feeble flash followed by a contraction of the gas indicates that 
combination has occurred. Finally, raise A until the mercury in it stands 
at the same height as before the explosion. The steam in A is now under 
the same pressure and at the same temperature as the mixed gases, and 
occupies two thirds the volume of the mixed gases taken. 

Hydrogen Dioxide, H 2 0, or H-O-O-H, is obtained by the 
slow evaporation of its aqneous solution as a colorless 
liquid, having a density of 1.45. It is soluble in all propor- 
tions in water, and is readily soluble in ether. The concen- 
trated aqueous solution gives off oxygen slowly at 20° and 
with explosive violence at 100°, hydrogen dioxide decompos- 
ing into water and oxygen: 

HA = H a O + 0. 

Dilute solutions are more stable, and when distilled a small 
portion of hydrogen dioxide passes over with the steam. 

Hydrogen dioxide acts as a bleaching agent, though not as 
rapidly as chlorine. It is found useful in restoring oil paint- 
ings which have been blackened by lead sulphide, and has 
been used in medicine. It serves as a reagent in qualitative 
analysis, and for the preparation of a number of compounds. 
Hydrogen dioxide is a strong oxidizing agent, readily giving up 
half its oxygen to many bodies, and is remarkable for its reac- 
tion with certain metallic oxides, which are reduced by it to 
the metallic state. Thus if silver oxide, Ag o 0, is placed in 
a dilute aqueous solution of hydrogen dioxide, oxygen gas 
will escape, and finely divided metallic silver will remain: 

Ag,0 + HA = 2Ag + H,0 + O s . 



120 THE SIXTH GROUP. 

Ozone in contact with a solution of hydrogen dioxide is 
changed to ordinary oxygen, thus : 

O s + H g O, = H,0 +20,. 

Hydrogen dioxide sets free iodine from metallic iodides, and 
the following reactions are employed in testing for the pres- 
ence of it in solutions : 



2KI + 


H 2 2 = 2KOH + I 2 . 


Potassium iodide 


Potassium hydroxide 


2FeI 2 + 


3H.O, = 2Fe(0H), + 2I 2 . 


Ferrous iodide 


Ferric hydroxide 



The first reaction takes place very slowly in highly dilute 
solutions, and immediately in concentrated solutions. But 
little iodine will remain in the free state, as the potassium 
hydroxide will combine with it. The solution, however, 
will react alkaline, and also for iodine with starch. Fer- 
rous iodide is decomposed immediately by very dilute solu- 
tions of hydrogen dioxide. 

Compounds containing the group OH are called hydroxides, 
and to OH the name Hydroxyl has been given. Hydrogen 
dioxide is free hydroxyl. Its molecular weight has not been 
determined, but the molecule H 2 2 is assumed because of its 
decomposition into water and oxygen. 

Water formed by synthesis, as in Exp. 76, does not contain 
a trace of hydrogen dioxide. But when a hydrogen flame 
impinges on water for a short time hydrogen dioxide may 
easily be detected, making evident the formation of it in the 
union of hydrogen and oxygen in the presence of water. 
The simplest view of the reaction is that a molecule of hydro- 
gen, H 2 , unites with a molecule of oxygen, 2 , to form H 2 2 . 

If it is true that hydrogen dioxide is the first product of 
the combustion of hydrogen with oxygen, then the formation 
of water may be due to the decomposition of hydrogen diox- 
ide by heat, thus : 

H,0. = H,0 + 0. 



HYDROGEN DIOXIDE. 121 

The free atom of oxygen may unite directly with a hydro- 
gen molecule to form another water molecule : 

H, + = H 2 0. 

Or the free oxygen atom may unite with another free oxygen 
atom to form 2 . 

Traube, who discovered this synthesis of hydrogen dioxide, 
suggests that the reaction may be as follows : 

H 2 + 0,=2H,0 + H s O a . 

(1) H 2 + H 2 

(2) H 8 2 + H, = 2H 2 0. 

The second equation represents the formation of water by 
the reducing action of hydrogen on hydrogen dioxide. 
That the formation of hydrogen dioxide is not due to the ad- 
dition of oxygen to the water molecule is supported by the fact 
that the most powerful oxidizing agents are without action on 
water. 

It is possible that the presence of water is essential to the 
combination of hydrogen and oxygen in the same way that it 
is in the reaction between carbon monoxide and oxygen, which 
will be considered later. The absence of hydrogen dioxide 
in the ordinary synthesis of water is owing to its instability 
at high temperatures. When hydrogen burns on the sur- 
face of water some of the hydrogen dioxide dissolves in the 
cool water, and is thus removed from the action of heat. In 
the combustion of carbon monoxide and also of illuminating 
gas on the surface of water hydrogen dioxide is formed. 

Hydrogen dioxide is commonly prepared by acting on ba- 
rium dioxide, Ba0 2 , with a dilute acid: 

Ba0 2 + 2HC1 = HA + BaCl a . 

The barium chloride, BaCl a , is a soluble salt, and remains 



122 THE SIXTH GROUP. 

in solution. "When sulphuric acid is used, barium sulphate, 
BaS0 4 , is precipitated, and may be removed by nitration: 

Ba0 2 + H 2 S0 4 = H 2 2 + BaS0 4 . 

The dilute solution may be concentrated by evaporation in 
a vacuum over sulphuric acid. The presence of a little free 
acid renders it less liable to decompose. 

Exp. 79.— Dilute hydrochloric in a test-tube with three or four times 
its bulk of water, and add gradually barium dioxide until it ceases to 
dissolve readily. The cheap barium dioxide of commerce will answer, 
although it is too impure for the preparation of concentrated solutions. 
Add to a small portion of the solution a few drops of thin starch-paste 
containing a little potassium iodide. 

It has already been shown that other substances (what ?) set free 
iodine from potassium iodide; hence the reaction is not characteristic for 
hydrogen peroxide. 

Exp. 80. — Dilute a portion of the hydrogen dioxide solution until no 
coloration appears on adding starch-paste and potassium iodide. Then 
add one drop of a dilute solution of ferrous sulphate or a minute frag- 
ment of the salt to the solution. The blue color will be seen at once. 
The ferrous sulphate, FeS0 4 , and potassium iodide react with formation 
of ferrous iodide and potassium sulphate, thus : 

FeS0 4 + 2KI = Fel 2 + K 2 S0 4 . 

As already stated, hydrogen dioxide sets free iodine from ferrous 
iodide. An excess of ferrous sulphate interferes with the reaction. 
Dilute solutions of hydrogen dioxide containing starch and potassium 
iodide become blue on standing some time. 

Exp. 81. — Hold a small jet of hydrogen burning at the end of a glass 
tube 3 or 4 mm. wide on the surface of water in a small porcelain dish. 
After some minutes test the water for hydrogen dioxide. The hydrogen 
for this experiment is best obtained from dilute sulphuric acid and zinc. 
That from impure hydrochloric acid does not answer. For the three 
following tests the dilute solution of hydrogen peroxide furnished by 
druggists answers. 

Exp 82. — To a solution of hydrogen peroxide, which should be nearly 
neutral, add a solution of pure potassium iodide. Free iodine will be 



CHLORINE MONOXIDE — HYPOCHLOROUS ACID. 123 

evident from the red color, and the solution will react alkaline to deli- 
cate red litmus paper, owing to the presence of potassium hydroxide. 

Exp. 83.— Place a test-tube containing not too dilute a solution of hy- 
drogen dioxide in hot water; oxygen gas will escape. 

Exp. 84. — Add silver oxide to a solution of hydrogen dioxide. If the 
solution is not too dilute, sufficient oxygen may be obtained to test. 



The Oxides and Hydroxides of the Seventh Group.* 

Chlorine Monoxide, Hypochlorous Oxide, C1 2 or C1-0-C1, 

is a yellowish-colored gas with a penetrating odor, resembling 
that of chlorine. Its observed gas density corresponds to 
that required by the formula C1 2 0. It condenses to a liquid, 
boiling at about — 19 °, which explodes violently when poured 
from one vessel to another, or when heated. Chlorine mon- 
oxide is prepared by passing dry chlorine over dry mercuric 
oxide, which reacts with formation of mercuric chloride and 
hypochlorous oxide : 

HgO + 2Cl 2 = HgCl 2 + Cl 2 0. 

Hypochlorous Acid, HC10 or C1-0H, is known only in di- 
lute aqueous solution. It is formed when hypochlorous oxide 
is dissolved in water. It is prepared by adding precipitated 
mercuric oxide to chlorine water: 

HgO + 2C1 2 + H 2 = HgCl 2 + 2C10IL 

On distillation the dilute acid passes over, and mercuric 
chloride remains. Hypochlorous acid can also be obtained 
from its salts. It is a powerful bleaching agent, yielding its 



* It is better to become familiar with the theory of acids, bases, and 
salts, and with the hydroxides of sulphur, before studying the oxides and 
hydroxides of this group. 



124 OXIDES A^D HYDROXIDES OF SEVENTH GROUP. 

oxygen readily to other substances with formation of hydro- 
chloric acid, thus : 

C10H = HC1 + 0. 

The most important salt of hypochlorous acid is calcium 
hypochlorite, which is continued in bleaching-powder. 

Chlorous Acid, HC10 2 . — This acid has not been isolated. 
"When a solution of chlorine peroxide is treated with a base, 
a mixture of chlorite and chlorate is formed : 

2C10 2 + 2K0H = KC10 2 + KC10 3 -f H 2 0. 

The chlorites are unstable compounds. Chlorous anhydride, 
C1 2 3 , has not been obtained. What was formerly regarded 
to be this compound has been shown to be a mixture of chlo- 
rine peroxide and chlorine. 

Chlorine Peroxide, C10 2 , is a dark-yellow gas which con- 
denses to red liquid, boiling at 9°. It is obtained by adding 
potassium chlorate to oil of vitriol. In the reaction chloric 
acid is formed, and is decomposed at once into perchloric acid 
and chlorine peroxide: 

KC10 3 + H 2 S0 4 = HC10 3 + HKS0 4 , 
3HC10 3 = HC10 4 + 2C10 2 + H 2 0. 

The preparation requires care, as the peroxide is explosive. 
The observed gas density of chlorine peroxide corresponds to 
the formula C10 2 . The following experiment illustrates the 
oxidizing action of the peroxide, and also the color of its solu- 
tion : 

Exp. 85. — Drop into a conical wine-glass filled with water a few 
crystals of potassium chlorate and a small bit of phosphorus. Pour 
some oil of vitriol, by means of a funnel tube, upon the chlorate The 
phosphorus will burn under water, taking oxygen from the chlorine 
peroxide set free. 



CHLORIC ACID— POTASSIUM CHLORATE. 125 

Chloric Acid, HC10 3 or C10 2 -0H, is best prepared by acting 
on barium chlorate with the requisite quantitiy of dilute sul- 
phuric acid: 

Ba(C10 3 ) 2 + H 2 S0 4 = BaS0 4 + 2HC10 3 . 

The barium sulphate is allowed to settle, and the clear solution 
is decanted and allowed to evaporate in vacuum over oil of 
vitriol. The residue remaining contains 40 per cent of HC10 S , 
and cannot be further evaporated without decomposition. 
The concentrated aqueous solution of chloric acid is colorless, 
has a strong acid reaction, and oxidizes paper, wood, and 
other organic substances. 

Potassium Chlorate, KC10 3 or C10 2 -0K.— This salt is the 
most important of the chlorates. It is used in the manu- 
facture of matches, in medicine, in explosive mixtures, and 
for other purposes. It is manufactured by passing chlorine 
gas into a hot mixture of lime, potassium chloride, and water: 

KC1 + 3CaO -j- 3C1 2 = KC10 3 + 3CaCl 2 . 

The chlorate is separated from the very soluble calcium 
chloride by crystallization. 

When a solution of potassium hydroxide is saturated with 
chlorine, potassium hypochlorite and chloride are formed, 
thus: 

2KOH + 01 2 = KCIO + KC1 + H 2 0. 

If the solution is left for a day, or is boiled, the hypochlorite 
decomposes into chlorate and chloride: 

3KC10 = KC10 3 + 2KC1. 

Potassium chlorate forms tabular crystals which arc per- 
manent in air. It is very soluble in hot water; but only 3.3 
parts of the salt are soluble in 100 parts of water at 0°. The 
salt is decomposed by heat with evolution of oxygen: the 



126 OXIDES AXD HYDROXIDES OF SEVENTH GROUP. 

chemical changes involved in the reaction are given under 
potassium perchlorate. 

Potassium chlorate yields its oxygen to combustible bodies 
more readily than potassium nitrate, and mixtures of it with 
sulphur, nitrobenzene (C 6 H 5 X0 2 ), or other substances explode 
with greater violence than gunpowder. Some of the mixtures 
containing chlorate are exploded by friction or percussion, 
and are used to fire gunpowder and other explosives. 

Exp. 86. — Saturate with chlorine a solution made by dissolving 50 
grams of potassium hydroxide in 150 cc. of water, and then boil a few 
minutes. Allow the solution to cool slowly ; crystals of potassium 
chlorate will be deposited. Separate them from the mother-liquor, 
rinse once or twice with cold water ; recrystallize by dissolving in a 
little more boiling water than required for solution, cool, and remove 
the crystals as before. Test them for potassium chloride by dissolving 
in water and adding a solution of silver nitrate. If the chlorate is free 
from chloride no precipitate will appear, as silver chlorate is soluble. 
Test the mother-liquor from the first crop of crystals for chlorine. 

Exp. 87. — Mix cautiously on a paper equal bulks of powdered potas- 
sium chlorate and sugar. Place the mixture on a brick, under a hood, 
and drop upon it a little oil of vitriol. 

Exp. 88. — Mix sulphur (flowers) with twice its bulk of pulverized 
potassium chlorate. Place as much of the mixture as can be held on 
the end of a penknife blade on an anvil, and strike with a hammer. 

Perchloric Acid, HC10 4 or C10 3 -0H.— This acid is best pre- 
pared by distilling pure potassium perchlorate with four times 
its weight of concentrated sulphuric acid. At 110° the pure 
acid, HC10 4 , distils over and condenses to a colorless or 
slightly yellow liquid. On continuing the distillation, the 
liquid which passes over changes to a crystalline mass of the 
hydrate, HC10 4 -f H 2 0, due to a partial decomposition of per- 
chloric acid into oxides of chlorine and water, the latter 
uniting with the pure acid in the distillate. 

Perchloric acid is a fuming, hygroscopic liquid, which 
explodes spontaneously when kept some days even in the 
dark. It decomposes with great violence when dropped upon 



POTASSIUM PERCHLORATE — BROMIC ACID. 127 

wood and charcoal. It forms with water crystals of the 
hydrate HC10 4 + H 2 0, melting at 50°. Perchloric acid is a 
strong monobasic acid. Its salts, the perchlorates, require a 
higher temperature for their decomposition than the chlorates, 
and are, moreover, not decomposed by hydrochloric acid. 

Potassium Perchlorate, KC10 4 or C10 3 -0-K. — When potas- 
sium chlorate is fused oxygen is evolved, and after a time the 
mass becomes pasty, and consists of potassium perchlorate, 
potassium chloride, with commonly some potassium chlorate. 
At a higher temperature the perchlorate is resolved into 
potassium chloride and oxygen. In order to prepare the per- 
chlorate, the mixture obtained by cautiously heating potassium 
chlorate is allowed to cool, pulverized, and the greater part of 
the chloride dissolved out by water. The residue is then 
heated with hydrochloric acid, which converts any remaining 
chlorate, with evolution of chlorine and oxides of chlorine, 
into potassium chloride, which is finally removed by washing 
with water. Potassium perchlorate dissolves in 65 parts of 
water at 15°; and in much less boiling water. 

Hypobromous Acid, HBrO, is similar in properties to its 
analogue, hypochlorous acid. 

Bromic Acid, HBr0 3 or Br0 2 -0H, is formed by the reaction 
of chlorine on bromine water, thus: 

Br, + 5C1 2 + 6H 2 = 2HBrO s + 10HC1. 

The acid is, however, best obtained by decomposing bromates. 
It is known only in aqueous solution, which is colorless, and 
which decomposes at 100°, with evolution of oxygen and 
bromine. 

Iodine Trioxide, I 2 3 , is said to be formed, together with the pentox.- 

ide, by the action of ozone upon iodine. 



128 OXIDES AND HYDROXIDES OF SEVENTH GROUP. 






Iodine Pentoxide, Iodic Anhydride, I 2 5 or I0 2 -0-I0 2 , 

is obtained as a white powder by heating iodic acid to 200°. 
It dissolves in water with formation of iodic acid. 

Iodic Acid, HI0 3 , or I0 2 -0H, is prepared by dissolving 
iodine in fuming nitric acid, evaporating the solution, and 
heating the residue until all nitric acid is expelled. The 
product so obtained is dissolved in water, and from the solu- 
tion rhombic crystals of the acid separate on evaporation. 

Iodic acid is regarded by some writers as a bibasic acid, 
having the formula H 2 I 2 6 , since it forms normal and acid 
salts, as for example, K 2 I 2 6 , HKI 2 6 . Sulphur dioxide and 
hydrogen sulphide reduce iodic acid, with separation of iodine. 

Periodic Acid, HI0 4 , has not been isolated, but its salts 
have been prepared. The hydrate HI0 4 -f- 2H 2 0, which may 
be viewed as monometaperiodic acid, IO(OH) 5 , is formed by 
the action of iodine on perchloric acid: 

CIO3-OH + I + 2H 2 = IO(OH), + 01. 

The acid is also formed by decomposing silver periodate with 
bromine. Monometaperiodic acid forms colorless deliques- 
cent crystals, melting at 133°, and decomposing at 140° into 
iodic anhydride, water, and oxygen. Periodic acid is remark- 
able for its complex salts, of which the following are ex- 
amples: Ag 3 I0 5 , Ag 3 H 2 I0 6 , Ag 4 I 2 9 , and Ba 6 I 2 12 . 

Manganese Monoxide, Manganous Oxide, MnO or Mn = 0, 

may be obtained by heating manganous carbonate in hydrogen. 
It is a grayish-green powder, which readily oxidizes when 
heated in air. 

Manganous Hydroxide, Mn(0H) 2 , forms a white, bulky 
precipitate on adding an alkali hydroxide to a solution of 
manganous chloride. It rapidly absorbs oxygen from the air, 



POTASSIUM MANGANATE. 129 

and turns brown. It is a basic hydroxide, forming manganous 
salts with acids. 

Manganous-Manganic Oxide, Mn 3 4 , occurs as the mineral 
hausmanite. This oxide is formed when any of the manga- 
nese oxides or manganous carbonate is strongly ignited in the 
air. 

Manganic Oxide, or Manganese Sesquioxide, Mn 2 3 , occurs 
as braunite. The hydrate, or hydroxide, Mn 2 3 -j- H 2 = 
2MnO-OH, is the mineral manganite. 

Manganese Dioxide, Mn0 2 , occurs as the mineral pyrolu- 
site, and is commonly known as black oxide of manganese. 
It is used in the preparation of chlorine, and in the manu- 
facture of glass to neutralize the color which iron imparts. 
Manganese dioxide is obtained artificially by igniting man- 
ganous nitrate. It separates as a hydrate on addition of a 
hypochlorite to a solution of manganese, and also when a dilute 
solution of manganese containing sodium acetate and a little 
free acetic acid is treated with chlorine or bromine. This last 
reaction is applied in the separation of manganese in chemical 
analysis. 

0— K 

Potassium Manganate, K 2 Mn0 4 or MnO n < K . — When 

an oxide of manganese is fused with potassium hydroxide in 
contact with air, or better, mixed with potassium chlorate, a 
dark-green mass results. This forms with water a green 
solution, which on evaporation in vacuum deposits small 
crystals of potassium manganate, isomorphous with potassium 
sulphate. A better method for the preparation of the man- 
ganate is to boil a solution of potassium permanganate ami 
potassium hydroxide as long as oxygen is evolved. The solu- 
tion on cooling yields a powder which is dissolved in potassium 
hydroxide, and from this last solution almost black crystals 
9 



130 OXIDES AND HYDROXIDES OF SEVENTH GROUP. 

of the manganate separate. Potassium manganate is decom- 
posed by water, thus: 

3MnO a <QJ| + 2H 2 = 2MnO,-0-K + 4X-OH + Mn0 2 . 

Carbonic acid facilitates the change by combining with the 
potassium hydroxide to form carbonate. In presence of suf- 
ficient alkali the decomposition does not occur. 

Exp. 89. — Fuse on platinum foil sodium carbonate to which a very 
small quantity of an oxide of manganese has been added. Sodium 
manganate will be formed, and will color the mass a bright green. 
This is a delicate test for manganese. 

Exp. 90. — To a concentrated solution of potassium hydroxide add a 
small crystal of potassium permanganate. The color of the solution of 
permanganate will change on warming to green, due to formation of 
manganate. Add more permanganate, and boil gently until the con- 
version into manganate is complete, shown by the pure green color of 
the solution, when some potassium manganate will separate. Pour the 
product into water. The dilute green solution, owing to the presence 
of excess of alkali, is tolerably permanent, but if sufficiently dilute will 
slowly change to red, owing to formation of permanganate. To the 
green solution of potassium manganate add an excess of sulphuric acid, 
and note change. 

Potassium Permanganate, KMn0 4 or Mn0 3 -0-K, is ob- 
tained by decomposing potassium manganate as already de- 
scribed. It forms dark-red crystals, which are isomorphous 
with potassium perchlorate, the two salts crystallizing together 
in all proportions. Potassium permanganate is soluble in 
about 16 parts of water at ordinary temperature. The solu- 
tion, as well as the dry salt, is a powerful oxidizing agent, and 
on this account is employed as a disinfectant. 

Ex. 91. — a. Dissolve a small quantity of oxalic acid in dilute sul- 
phuric acid, and add drop by drop a solution of potassium perman- 
ganate until the red color of the permanganate is permanent. The 
reaction is represented as follows . 



VALENCE OF ELEMENTS OF SEVENTH GROUP. 131 

5 600H+ 2Mn ° 3 -°- K + 3S ° 2 <OH = 

S ° 2 <0-K + 2S ° 2 < 8 > Mn + 10C ° 2 + 8H2 °- 

b. Add to river or spring water sufficient potassium permanganate to 
impart a reddish tint. After some days the color will disappear, owing 
to reduction of the permanganate by the organic matter in the water. 

c. Leave a piece of paper for several days in a solution of potassium 
permanganate, and note change. 

If dry potassium permanganate is added to well-cooled sulphuric acid, 
the latter is colored green, and, on adding more permanganate, Man- 
ganese Heptoxide, Mn 2 7 , separates as a heavy liquid, with a metallic, 
greenish lustre. Manganese heptoxide explodes with incandescence 
when heated. It forms a green solution with concentrated sulphuric 
acid, and dissolves in water to a red solution, probably of permanganic 
acid. When sulphuric acid containing an excess of potassium perman- 
ganate is cautiously heated to 50°, Manganese Trioxide, Mn0 3 , is evolved 
as a violet fume, which may be condensed to a dark-red liquid. It de- 
composes without incandescence when heated into manganese dioxide 
and oxygen. It forms an unstable red solution with water, which is 
supposed to contain manganic acid, H 2 Mn0 4 . 

Permanganic Acid, HMn0 4 or Mn0 3 -0H, is prepared by- 
mixing barium permanganate with the required amount of 
sulphuric acid. The deep-red solution thus obtained decom- 
poses on standing, quickly when heated, with evolution of 
oxygen and separation of hydrous manganese dioxide. 

Permanganic Oxychloride, Mn0 3 Cl, is known. It is the 
chloranhydride of permanganic acid. 



Yalence of the Elements of the Seventh Group, and 
Constitution of their Oxygen Com pounds. 

The halogens are univalent towards hydrogen and metals ; 
in their oxygen compounds they exhibit higher valence. 



132 OXIDES AXD HYDROXIDES OF SEVENTH GROUP. 

Since all views of the valence of an element are based on the 
number of atoms in the molecules of its compounds, it is 
obvious that the hypothesis of valence can be strictly applied 
only to compounds whose molecular weights are known. 

It is evident that if the number of atoms in a molecule of 
a given compound are unknown the valence assigned to the 
atoms becomes more or less hypothetical. Nevertheless it 
should be remembered that a constitutional formula in which 
the valence of atoms or groups of atoms (radicals) is implied 
or marked, thereby indicating its constitution or structure, 
represents more of the chemical properties of a compound 
than an empirical formula, that is, one which represents only 
the relative number of atoms in the compound. 

Kegarding valence as the capacity of an atom to hold other 
atoms in combination, we must assume that the valence of 
iodine in I 2 0. is greater than in HI. Later we shall see that 
valence measured by oxygen may be higher than valence 
measured by hydrogen. 

Oxygen acids are generally assumed to contain the radical 
hydroxyl, OH, that is, to be hydroxides. Proof of this will 
be given in the study of the constitution of sulphuric acid 
and organic acids. Viewing the oxy-acids of chlorine as hy- 
droxides, we have the formulas 

Cl-OH, (CIO)-OH, and (C10J-OH. 



Chlorous acid may be represented by Cl-O-OH or 
in 
= Cl-OH, according as we assume the chlorine in it to be 

a monad or a triad. In the first formula the radical hvdroxvl 

is joined to oxygen, and in the second to chlorine. There 

are no reactions of the acid which support the constitution of 

either formula, but it is assumed that the hydroxyl is linked 

to chlorine, since, in other acids, for example those containing 

carbon, hydroxyl is not joined to oxygen. The argument 



VALENCE OF ELEMENTS OF SEVENTH GROUP. 



133 



will be given later in the discussion of the constitution of 
carbon compounds. 

Chlorine peroxide forms with water chlorous and chloric 
acids, and is therefore regarded as the mixed anhydride of 
these acids, thus : 







in v o 

C1-0-C1^q 



H n O 



Chlorous chloric anhydride 



III f) V 

= Cl-OH + ^Cl-OH. 

Chlorous acid Chloric acid 



Bromic and iodic acids are probably analogous in constitu- 
tion to chloric acid. 

Perchloric acid crystallizes with one molecule of water, and 







OH 



the compound may oe a trihydroxide, thus: ^^Cl^-OH. 

U M)H 



This constitution is doubtful, since no salts of a tribasic per- 
chloric acid are known. The monobasic character of the acid 



is represented by the formula CI 




Orthoperiodic acid, I(OH) 7 , and its salts are unknown. 

VII 

Monometaperiodic acid, OI(OH) 5 , may be viewed as derived 
from the ortho-acid by the removal of the elements of one 
molecule of water, thus: 





Orthoperiodic acid 



Monometaperiodic add 



134 



OXIDES AXD HYDROXIDES OF SEYEXTBZ GROUP. 



Dimetaperiodic acid and trinietaperiodic acid have not been 
isolated, but their salts have been prepared. The diineta- 
acid may be regarded as derived from the monometa-acid, 
and the trimeta- from the dimeta-acid, thus: 




vn 
H„0=I 




VII 

H„0 = I 




Monometaperiodic 
acid 



Dimetaperiodic 
acid 



Trimetaperiodic 
acid 



It has been stated that potassium perchlorate and potassium 
permanganate are isomorphous. It is therefore assumed that 
perchloric and permanganic acids have an analogous constitu- 
tion. The anhydride of permanganic acid is Mn 2 7 , whose 



structural formula is 








0, on the assumption that 









each oxygen atom is linked by two bonds to manganese. 



Permanganic acid is Mn 




The structural formulas given of the oxy-acids of the 
seventh group are analogous to the formulas of ortho- and 
meta-acids of other elements, which will be studied later. 
Even if the foregoing formulas be considered as purely hypo- 
thetical, they will be found useful as an aid to the memory. 



SUMMAKY OF THE SEVENTH GBOUP. 



135 



The student lias only to recollect that an ortho-acid of a 
heptavalent element has seven hydroxyls, and that the meta- 
acids are derived from it by the withdrawal of the elements 
of one or more molecules of water. 

The members of the seventh group exhibit in their differ- 
ent classes of compounds valence from one to seven, thus: 



Fluorine, . . 


, I 














Chlorine, . . 


. I 




in 




V 




VII 


Bromine, . . 


. I 




in 




V 






Iodine, . . . 


. I 








V 




VII 


Manganese, 




II 




IV 




VI 


VII 



Summary of the Seventh Group. 



It has been stated in the summary of the halogens that they 
exhibit decreasing affinity for hydrogen with increasing atomic 
weights; towards oxygen the reverse is true, viz., increasing 
affinity, for oxygen with increasing atomic weights. Fluorine, 
with the lowest atomic weight in the group, does not unite 
with oxygen; and fluorine hydroxides are unknown. 

In this connection it is to be further noticed, that oxides of 
iodine are more stable than those of chlorine. It is also 
remarkable that the stability of the oxy -acids of the halogens 
increases with the addition of oxygen. Thus hypochlorons 
acid, ClOH, is a weak acid, being liberated from its salts by 
carbonic acid; while perchloric acid forms stable salts, from 
which it is set free only by a strong acid, such as sulphuric. 
The following are the oxides and hydroxides of the seventh 
group : 



136 OXIDES AXD HYDEOXIDES OF SEVEXTH GKOUP. 



C1 2 0. 



C10H. 



BrOH. 



Hypo- Hypo- No Ox- Hypo- 
chlorous chlorous ides of bromous 
anhydride acid Bromine acid 



CIO. OH. 




I 2 a 



Chlorous 
acid 



MnO. MnCOH), 



Manganous Manganous 
oxide hydroxide 

Mn 2 3 . MnO. OH. 

Manganic Manganic 
oxide hydroxide 



Ci0 2 . 

Chlorine 
Peroxide 



C10 2 .0H. 

Chloric 
acid 



Br0 2 .OH. I 2 5 . I0 2 .0H. 

Bromic Iodic Iodic 

acid anhy- acid 

dride 



C10 3 OH. 

Perchloric 
acid 



I0 3 OH. 

Periodic 
acid 



MnO, 



Manganese 
dioxide 



M11O3. Mn0 2 (OH) 2 . 

Manganic Manganic 
anhydride acid 

Mn 2 07. M11O3.OH. 
Perman- Perman- 
ganic ganic acid 



anhydride 



Chlorine, bromine, and iodine are closely related in chemi- 
cal properties, and form a well-defined sub-group. Fluorine, 
with the lowest atomic weight in the group, stands apart from 
the other members in that it forms no oxides and hydroxides. 
The hydrogen compounds and the hydroxides of the halogens 
are acids, and do not possess any basic characteristics. Man- 
ganese, because of the magnitude of its atomic weight, is 
classed in the seventh group. It shows, however, relation- 
ship to the halogens only in permanganic acid and its salts. 
The lower oxides and hydroxides of manganese are basic com- 
pounds, forming salts with acids. The element in its basic 
character and metallic properties resembles iron. Manganese 
dioxide is an indifferent oxide, neither basic nor acidic. 
Manganic acid is analogous to sulphuric acid. The seventh 
group, therefore, presents marked analogies and great differ- 
ences in properties. 



OXIDES AND HYDROXIDES OF ALKALI METALS. 137 

The Oxides and Hydroxides of tlie First Group. 

The hydroxides of the alkali metals are remarkable for their 
strongly basic properties and for their stability at high temper- 
atures, not being decomposed into water and oxide even at a 
red heat. They constitute the caustic alkalies, so called from 
their caustic or corrosive properties. The oxides of copper 
and silver are more permanent than their hydroxides, and are 
strong basic oxides; but the hydroxides of gold are acids. 
The oxides of copper are not decomposed by ignition, while 
the oxides of silver and gold are reduced to the metals when 
heated. 



The Oxides and Hydroxides of the Alkali Metals. 

Lithium Oxide, Li 2 or Li-0-Li, is formed together with a higher 
oxide when lithium burns in the air It may be prepared by igniting 
lithium nitrate. It is a white mass, which dissolves slowly in water with 
formation of the hydroxide. 

Lithium Hydroxide, Li-OH, resembles the other alkali hydroxides, but 
does not absorb carbon dioxide as rapidly, and, unlike potassium hydrox- 
ide, does not deliquesce in the air. 

Sodium Monoxide, Na 2 or Na-O-Na, is best prepared by 
heating sodium hydroxide with sodium : 

Na-OH + Na = Na-O-Na + H. 

It unites energetically with water to form sodium hydroxide. 

Sodium Dioxide, Na a 2 , or Na-0-O-Na. — When sodium is 
burned in air both the monoxide and dioxide are formed. 
The dioxide is also obtained by heating sodium in oxygon 
until the mass ceases to increase in weight. If added gradu- 
ally to water, sodium dioxide dissolves, and the solution 
on spontaneous evaporation yields crystals of Na.,0, + SI 1,0. 
When added to a small quantity of water, the mixture be- 



138 OXIDES AXD HYDROXIDES OF EIRST GROUP. 

comes hot, oxygen gas is evolved, and sodium hydroxide is 
formed. 

Exp. 92. — Burn a piece of sodium the size of a pea on a porcelain cru- 
cible cover which is heated with a small lamp flame. When cool, place 
cover and contents in about 100 cc. of water; acidify the solution with 
hydrochloric acid, and test for hydrogen dioxide. 

Express by an equation the reaction between the sodium dioxide and 
acid whereby hydrogen dioxide results. 

Sodium Hydroxide, Sodium Hydrate, Caustic Soda, NaOH 
or Na-OH. — The cheap and common method of making 
caustic soda is by treating a boiling solution of sodium car- 
bonate with slaked lime: 

CO< 0-la + Ca< OH = 2Xa -° H + CO<3>Ca. 

Sodium carbonate Calcium hydroxide Sodium hydroxide Calcium carbonate 

The insoluble calcium carbonate may be removed by fil- 
tration or be allowed to subside, and the clear liquid drawn 
off with a siphon. The solution is evaporated in an iron ves- 
sel, the heat being finally raised to a dull red, and the molten 
mass is cast into sticks or other convenient forms. 

Pure sodium hydroxide is best obtained by dissolving metal- 
lic sodium in water, and evaporating the solution in a silver 
dish. 

Sodium hydroxide is an opaque, brittle, white solid. It de- 
liquesces in moist air, absorbs carbon dioxide, C0 2 , and in time 
is completely changed to sodium carbonate, which remains as a 
white powder. When a solution of sodium hydroxide of den- 
sity 1.385 is cooled below 0°, transparent crystals of the hy- 
drate, 2XaOH -f- TH 2 0, separate. From alcoholic solutions 
the hydrate, XaOH -f 2H 2 0, has been obtained. 

Caustic soda is the cheapest of the alkalies, and is exten- 
sively used in the arts and in laboratories. Its aqueous solu- 
tion is known as soda lye. 



POTASSIUM HYDROXIDE. 139 

Exp. 93. — Slake 6 grams of quick-lime in 200 cc. of boiling water 
in a flask, and add 20 grams of crystallized sodium carbonate, 

CO<!^~iv a H - 10H 2 O. Boil some minutes, then "filter a portion, and 

to filtrate add an excess of hydrochloric acid. If gas (carbon dioxide) is 
given off the sodium carbonate has not all been converted into hydroxide, 
and a little more lime should be slaked in hot water, and the milk of lime 
added to the original portion and the whole boiled again. When a part 
of the clear solution does not evolve gas, on adding an excess of hydro- 
chloride acid filter the solution into a jar. Try the action of the sodium 
hydroxide solution on the fingers. 

To a solution of sodium carbonate add acid; the escaping gas is car- 
bon dioxide. Test the white substance which remained on the filter 
when the sodium hydroxide was filtered for carbon dioxide, thus: Rinse 
some of it into a test-tube with a jet of water from the wash-bottle, then 
add hydrochloric acid. 

Exp. 9i.— a. Leave a small piece of sodium hydroxide exposed to the air 
in a porcelain dish for some days. Observe changes, and finally test for 
carbon dioxide. 

b. Allow the solution of sodium hydroxide of the previous experiment 
to remain in an open jar several days, then test it for carbon dioxide. 

Potassium Oxides. — Potassium burns in dry air with for- 
mation of a mixture of two oxides: the monoxide, K 2 or 
K-O-K, and the tetroxide, K 2 4 or K-O-O-O-O-K. The 
monoxide resembles its sodium analogue in properties. The 
tetroxide has a dark-yellow color, and decomposes at a white 
heat into monoxide and oxygen. It dissolves in water with 
formation of potassium hydroxide, hydrogen dioxide, and free 
oxygen. 

Potassium Hydroxide, Potassium Hydrate, Caustic Potash, 
KOH or K-OH, is prepared from potassium carbonate, 

CO<q_-£-, by treatment with lime in the same manner as 

sodium hydroxide. 

Potassium hydroxide is,white and brittle, ami melts below a 
red heat. It absorbs water and carbon dioxide from (he air. 
deliquescing, and forming potassium carbonate, which also 



140 OXIDES AKD HYDROXIDES OF FIRST GROUP. 



deliquesces if the air is not too dry. It dissolves in about 
half its weight of water, with evolution of much heat. The 
hot saturated solution deposits on cooling colorless crystals 
of the hydrate, KOH -f- 2H 2 0, and the mother-liquor at 
ordinary temperature contains about 50 per cent of KOH. 
It is also soluble in alcohol. Solutions of potassium hydroxide 
are used to absorb carbon dioxide, and a saturated solution 
will remove water from gases. 

Potassium hydroxide is a powerful base, combining with 
acids to form salts, and is used to separate other hydroxides 
from solutions of salts, as we shall learn by subsequent experi- 
ments. The aqueous solution of potassium hydroxide is fre- 
quently called potash lye. It corrodes animal and vegetable 
matter, and when concentrated cannot be filtered through 
paper. It is used in making soft soap. 

Exp. 95. — Feel of a bit of caustic potash, and observe its corrosive 
action on the skin. Leave it in an open dish until it has deliquesced, 
then add to it a little water and an excess of hydrochloric acid. The 
escaping carbon dioxide was absorbed from the air. 

Rubidium Hydroxide, Rb-OH, and Caesium Hydroxide, 
Cs-OH, closely resemble potassium hydroxide. 






The Oxides and Hydroxides of Copper, Silver, and 

Gold. 

Cu-Cu 
Copper Tetrantoxide, Cu 4 or i > 0, is an olive green powder, 

Cu— Cu 

which absorbs oxygen from the air. 

Cu 
Cuprous Oxide, Cu o or i >0, also called the red oxide 
Cu 

of copper, may be obtained by heating finely divided copper 



CUPRIC OXIDE. 141 

and cupric oxide in the proper proportions, and also by 
warming an alkaline solution of copper to which glucose or 
other reducing sugar has been added. This reaction is applied 
in the estimations of sugars. 

Exp. 96. — To a few drops of a solution of cupric sulphate add a solu- 
tion of potassium hydroxide until the precipitate of cupric hydroxide 
which first forms dissolves. Add a very little glucose or a drop of 
molasses, and boil the solution. The red precipitate is cuprous oxide. 
A yellow precipitate of hydrated cuprous oxide sometimes results. 



Cupric Oxide, CuO or Cu = 0. — When copper is heated in 
air it oxidizes, and on plunging the red hot metal into water 
the coating comes off. This is known as copper scale, and is 
a mixture of cuprous and cupric oxide. It is completely con- 
verted into cupric oxide by prolonged heating to redness in 

air or oxygen. On igniting cupric nitrate, xrr) 9 Ir)>C'u, 

cupric oxide remains as a black powder, which sinters together 
on heating more strongly. Cupric oxide fuses at a high red 
heat, with loss of oxygen, and is partly converted into cuprous 
oxide. 

Hot copper oxide yields its oxygen readily to hydrogen and 
carbon, and is used in the combustion of compounds of these 
two elements in analysis. 

Exp. 97. — Heat a piece of copper to redness and thrust it into water. 
Repeat several times. The scale will show both black and red oxide. 

Exp 98. — Pass a very slow current of illuminating gas over glowing 
oxide of copper contained in a hard glass tube. The gas is a mixture 
of hydrogen, carbonic oxide, CO, and compounds of carbon and hy- 
drogen. These substances are burned in the tube, and the products are 
water and carbon dioxide, C0 2 , an invisible gas. 

OH 

Cupric Hydroxide, Cu(0H) 2 or Cu< QH . — When potassium 

hydroxide is added to a cold solution of a cupric salt a bulky 



142 OXIDES AXD HYDROXIDES OF FIEST GROUP. 

blue precipitate of cupric hydroxide is formed, which changes 
to cupric oxide on boiling the liquid. 

Exp. 99. — Make cupric hydroxide, and heat cautiously to boiling. 
Represent changes by equations. 

Copper Dioxide, Cu0 2 , is not known, but its hydrate, Cu0 2 -f- H 2 0, has 
been obtained. 

Silver Tetrantoxide or Argentous Oxide, Ag 4 0, and also some unstable 
argentous salts, have been prepared. 

Silver Oxide, Argentic Oxide, Ag 2 or Ag-O-Ag. — This 
oxide is prepared by precipitating silver nitrate with potas- 
sium hydroxide. If the solutions are cold the precipitate is 
brown, and on drying becomes darker, while from hot concen- 
trated solutions the oxide falls as a heavy black powder. 
Silver oxide decomposes completely below 300°, and yet is 
apparently formed to a slight extent when silver vapor cools 
in air. It is reduced by hydrogen at 100°. It is slightly 
soluble in water. The solution has a metallic taste and 
alkaline reaction, and may be assumed to contain silver hy- 
droxide, AgOH. Moist silver oxide absorbs carbon dioxide 
from the air. 

Exp. 100. — Dissolve 5 grams of silver nitrate in 100 cc. of water, add 
an excess of potassium hydroxide, wash the precipitate thoroughly on a 
filter, and dry it either at ordinary temperature or at 100°. Write the 
equation of the reaction. Heat a small portion of the oxide on the 
cover of a porcelain crucible. 

Silver Dioxide, Ag 2 2 , is formed in the electrolysis of a concentrated 
solution of silver nitrate, and, according to Schoenbein, by the action of 
ozone on silver. lies, by heating silver nitrate with silicic acid, has ob- 
tained a red modification, that made by electrolysis being black. 

Aurous Oxide, Au„0 or Au-O-Au, is formed when aurous 
chloride is treated with a cold dilute solution of potassium 
hydroxide. It does not yield aurous salts, but with hydro- 
chloric acid gives auric chloride and gold. 



OXIDES AND HYDROXIDES OF GOLD. 143 

Auric Hydroxide, Au(OH) 3 , is perhaps capable of separate 
existence ; but Schrottlaender, in attempting to prepare it by 
decomposing gold trichloride with potassium hydroxide or 
auryl nitrate with water, obtained products which, after long 
drying, corresponded in composition fairly well with the fol- 
lowing body: 

Auryl Hydroxide, AuO.OH, or = Au-OH. — This com- 
pound has been prepared by decomposing a dilute solution 
of auric chloride with sodium hydroxide, and then adding 
sodium sulphate, which separates the hydroxide as a dark- 
brown precipitate (Thomsen). Auryl hydroxide at 200° 
loses water, with formation of Auric Oxide, Au 2 3 , which 
loses all its oxygen at a higher temperature. The hydroxide 
dissolves in acids, yielding auryl salts, in which the basic 
radical may be regarded as -Au=0. 

Potassium Aurate, AuO.OK or = Au-0-K, is obtained m 
yellow crystals with three molecules of water, by dissolving 
auryl hydroxide in a slight excess of pure potassium hy- 
droxide, and evaporating the solution. The salt is not very 
stable, is easily soluble in water, and is used for gilding. 

The compounds of gold thus far considered show that the 
metal is basic or acidic in its nature, according as it is united 
with strongly acidic or basic elements. For example, it is 
basic in AuCl 3 , and acidic in AuO.OK. 



SULPHUR, SELENIUM, AND TELLURIUM. 

Sulphur, S. 

Atomic Wei glit, 32, Molecule, S 3 . 

Xatiye sulphur is abundant in many volcanic regions, and 
is often found in beautiful rhombic crystals. It occurs in 
combination with many metals, as in lead sulphide, PbS, 
and iron sulphide, FeS„, and is a constituent of many other 
minerals. Sulphur compounds are present in many animal 
and vegetable substances, and in natural waters. 

The process of extracting native sulphur is simple. The 
ore is heated until the sulphur melts and flows from the 
admixed earthy substances. The crude sulphur thus ob- 
tained is purified by distillation. If the vapor is conducted 
into a cool chamber, a fine yellow powder falls, which is known 
as flowers of sulphur. When the chamber becomes hot only 
liquid sulphur collects, and this is cast in wooden moulds, 
and is known as roll sulphur or brimstone. 

Sulphur exists in several allotropic forms. 

Rhombic Sulphur is found in transparent yellow octahedrons, 
variously modified. It is obtained by the slow evaporation of 
a solution of sulphur in carbon disalphide, and forms when 
sulphur vapor or a large mass of molten sulphur cools slowly. 
Rhombic sulphur has a density at 0° of 2.05, and melts at 
114°. 5, forming a yellow liquid having a density of 1.803. It 
is very soluble in carbon disulphide, but slightly soluble in 
alcohol and insoluble in water. 



SULPHUR. 145 

Exp. 101. — Dissolve sulphur in carbon disulpliide, and filter the solu- 
tion into a beaker. Cover the beaker with paper to keep out dust, and 
also to insure slow evaporation, and leave it under a hood or out of 
doors until the solvent has evaporated. 

Monoclinic Sulphur. — If molten sulphur is allowed to cool 
in the air until it has partly solidified and the liquid is then 
poured off, long delicate crystals will be seen, which are deeper- 
colored than the rhombic crystals of the foregoing experiment. 
This modification is not permanent ; and the crystals in a 
few days, and often in a few hours, become opaque, owing to 
a change to an aggregation of minute crystals of the rhombic 
form. Monoclinic sulphur melts at 120°, has a density of 
1.96, and is soluble in carbon disulpliide. The solution yields 
rhombic crystals when evaporated. 

Exp. 102. — Allow molten sulphur contained in a beaker to cool, watch 
the growth of the crystals, and when there has formed on the surface of 
the sulphur a crust, make a hole in it and pour the liquid portion into a 
porcelain dish. The crust may be cut away with a knife. Compare 
the crystals with those of the previous experimeut. Keep the sulphur 
in the beaker for a few days until it has become opaque. The portion 
in the porcelain dish when solid will have the color of the crystals, and 
in fact will be a mass of crystals, and will also undergo the same 
change. 

Amorphous Soluble Sulphur. — When an acid is added to a 
solution of an alkali polysulphide, sulphur separates in exceed- 
ingly fine particles. This form, which is known as milk of 
sulphur, is white, and is soluble in carbon disulpliide. Under 
the microscope it appears to be non-crystalline. 

Exp. 103. — Dissolve some flowers of sulphur in a hot concentrated 
solution of potassium hydroxide, dilute the solution with water, and add 
an excess of hydrochloric acid. If the liquid containing the milk of 
sulphur is poured upon a filter, some of the finely divided sulphur will 
pass through the pores of the paper. The composition of the sulphides 
of potassium will be given later. 



146 THE SIXTH GKOUP. 



Amorphous Insoluble Sulphur. — Sulphur separates in a yel 
low plastic amorphous form, insoluble in carbon disulphide, 
when a chloride of sulphur is decomposed by water or a solu- 
tion of a thiosulphate by hydrochloric acid. A plastic form, 
also insoluble in carbon disulphide, is obtained by quicklv 
cooling sulphur which has been heated to near its boiling 
point. Thus obtained it may be drawn into threads. It be- 
comes hard and brittle in a few days. When flowers of sul- 
phur are treated with carbon disulphide a part remains undis- 
solved. 

Sulphur at temperatures but little higher than its melting 
point is a pale-yellow liquid, which gradually becomes dark 
red and viscid on heating, and at 200° to 250° is too thick 
to pour readily. On increasing the temperature the sulphur 
grows less viscid, but remains dark-colored. Sulphur boils at 
448°. 4, under a pressure of 760 mm. The density of sulphur 
vapor diminishes with increasing temperature, but has been 
found constant between 860° and 1040°, and to be 32, cor- 
responding to the gas molecule S 2 . At 518° a density of 96 
has been observed, indicating the existence of molecules of 
S 6 . Blitz has, however, found that wheu sulphur is mixed 
with nitrogen, its density at 518° is from 65 to 71. This 
result, together with the fact that the density is not constant 
at moderate intervals of temperature between 500° and 600°, 
renders doubtful the existence of gas molecules of S 6 . 

Exp 104. — Melt clean pieces of roll sulphur in a test-tube, with care 
not to heat the sulphur much above its melting point, so as to obtain a 
thin yellow liquid, and pour a few drops into water. These will be 
brittle. Next heat the sulphur until it will not pour from the tube, then 
heat to boiling, and pour in a thin stream into water. Set aside a part 
of the plastic sulphur obtained, and after some days note any change. 






SELENIUM — TELLURIUM. 147 

Selenium, Se. 

Atomic Weight, 79. Molecule, Se 2 . 

This element occurs sparingly in nature. It is obtained in 
the manufacture of sulphuric acid from pyrites which contain 
traces of it. Selenium, like sulphur, exists in different modi- 
fications. When molten selenium is quickly cooled, an amor- 
phous, almost black mass results, soluble in carbon disulphide. 
From such a solution small dark monoclinic crystals, of density 
4. 5, separate on evaporation. When an aqueous solution of sele- 
nious acid is treated with sulphur dioxide, selenium separates 
as a red powder, which is also soluble in carbon disulphide. 
Metallic and insoluble selenium results when molten selenium 
is very slowly cooled. It has a lead-gray metallic lustre, and 
is a conductor of electricity. Its density is 4.5, and it melts 
at 217°. Minute black crystals of selenium, having a density 
of 4.8, separate when a strong solution of potassium selenide 
is exposed ta air. This black modification is also insoluble in 
carbon disulphide. 

Selenium boils below 700°. Its gas density at 1400° cor- 
responds to the molecule Se 2 . Selenium burns in air with a 
blue flame, and with a characteristic odor, said to resemble 
that of rotten horse-radish. 



Tellurium, Te. 

Atomic Weight, 125. Molecule, Te 2 . 

Tellurium is found in nature in small quantity, both free and 
in combination. It has a bluish-white metallic lustre, and a 
density of 6.24. It melts at about 500°, and may be distilled 



148 



THE SIXTH GROUP. 



in hydrogen gas. Its gas density at 1390° has been found to 
be 131.6. Tellurium is insoluble in carbon disulphide. 
When ignited in air it burns with formation of the dioxide. 



The elements oxygen, sulphur, selenium, and tellurium ex- 
hibit gradations in physical properties which are related to 
their atomic weights. Oxygen, which stands apart in chemical 
properties, likewise shows the greatest variation in physical 
properties, and possesses no metallic characteristics, while 
tellurium has the properties of a metal. 





Oxygen. 


Sulphur. 


Selenium. 


Tellurium. 


Atomic weight, 


16 


32 


79 


125 


Gas density, 


16-24 


32 


79 


125 


Density, 


1.12 liquid 


1.96-2.05 


4.5-4.8 


6.2 


Melting point, . 





114-120° 


217° 


500° 


Boiling point, . 


-181° 


448° 


below 700° 


Red heat 



Hydrogen Sulphide, Sulphuretted Hydrogen, H 2 S or H-F-H, 

is a colorless gas, occurring in many mineral waters com- 
monly called sulphur waters, to which it imparts medicinal 
properties. It is formed in the decay of various organic sub- 
stances containing sulphur, and to it is due in part the odor 
of rotten eggs. Hydrogen unites with boiling sulphur to 
form hydrogen sulphide, but the amount obtained is small. 
The common method of making hydrogen sulphide is to act on 
ferrous sulphide with dilute hydrochloric or sulphuric acid. 
The following equations represent the changes : 



Fe^S 



IHC1 



H 
H 



Fe=S + S0 2 <°g = S<g 



CI 
01. 

S0 2 <Q>Fe. 



The gas thus prepared contains a little hydrogen set free 
from the acid by particles of metallic iron in the ferrous sul- 



HYDROGEN SULPHIDE. 149 

phide ordinarily used. If native antimony trisulphide is em- 
ployed^ pure hydrogen sulphide will be obtained. 

The gas density of hydrogen sulphide is 17 ; its molecular 
weight is therefore 34. Analysis has shown that there are 
32 weights of sulphur and 2 weights of hydrogen in 34 weights 
of hydrogen sulphide. Hence its molecule is H 2 S. 

Hydrogen sulphide condenses to a liquid, boiling at —62° 
and solidifying at —85°. Water at common temperature dis- 
solves about three volumes of the gas. The solution reacts 
slightly acid, and has the odor and sweetish taste of the gas. 
Hydrogen sulphide is readily decomposed by heat and by 
oxidizing agents. Its aqueous solution gradually becomes tur- 
bid in contact with air owing to separation of sulphur and 
formation of water, thus : 

H 2 S + = H 2 0+ S. 

Nitric acid, chlorine, and bromine decompose hydrogen sul- 
phide. The reaction with chlorine is expressed by the 
equation — 

H 2 S + Cl 2 = S + 2HC1. 

Hydrogen sulphide is very inflammable, burning with a blue 
flame with formation of water and sulphur dioxide, SO a : 

H 2 S + 30 = H 2 + S0 2 . 

If the supply of air is not sufficient for complete combus- 
tion part of the sulphur separates. Hydrogen sulphide is 
much used in analytical chemistry for the precipitation o\' 
various metals as sulphides. Sulphuretted hydrogen is poi- 
sonous to inhale, and air containing a little of it produces 
nausea and headache. 

Exp. 105. — The Kipp apparatus, Fig. 74, is one of the most con- 
venient forms of generators for the preparation of hydrogen sulphide. 

The globe A contains ferrous sulphide. The aeid employed is common 



150 



THE SIXTH GROUP. 



muriatic acid, diluted with twice its bulk of water. The wash-bottle 
contains water to free the gas from any solution carried over. Instead 
of a Kipp generator, a bottle such as is used in Exp. 6, Fig. 52, will 
answer. 

Allow hydrogen sulphide to bubble slowly through distilled water in 
a bottle until, on closing the bottle with the thumb and shaking, the gas 
is no longer absorbed. Keep the solution for later experiments. 

Exp. 106 — a. Pass hydrogen sulphide through a hot tube. b. Fill a 
tall cylinder with the gas and ignite it. 




Fig. 74. 



Exp. 107. — a, Add to a solution of hydrogen sulphide chlorine or bro- 
mine water, and (b) to another portion its bulk of concentrated nitric acid. 

Exp. 108. — a. Treat a solution of cupric sulphate containing hydro- 
chloric acid with an excess of hydrogen sulphide. The following is 
the reaction: 



.0 



H 



S0 2 <~>Cu + S<S = Cu 



o 



11 



■s + so a <gg. 



b. Add hydrogen sulphide to a solution of zinc sulphate to which a 



CHLORIDES OF SULPHUR. 151 

little hydrochloric acid has been added. Note the result. Next add to 
the solution ammonia in excess. The white precipitate is zinc sulphide. 

c. Pass hydrogen sulphide into a solution containing copper sulphate, 
zinc sulphate, and some hydrochloric acid. After the copper is all pre- 
cipitated, filter, and add to the clear colorless filtrate ammonia to alkaline 
reaction. Some zinc sulphide will be precipitated, and more will 
separate on adding hydrogen sulphide. 

The separation of copper from zinc in the experiment illustrates the 
use of hydrogen sulphide in chemical analysis. 

Exp. 109. — a. Dissolve arsenious oxide in hot strong hydrochloric 
acid, but do not boil the mixture, as a volatile poisonous arsenic com- 
pound will be given off. Pour the solution into water, and add hydrogen 
sulphide. The precipitate is arsenic trisulphide. 

b. Dissolve tartar emetic, an antimony salt, in water, and add some 
tartaric acid and hydrochloric acid, and then hydrogen sulphide. The 
precipitate is antimony trisulphide. 

Hydrogen Disulphide, H 2 S 2 , is the analogue of hydrogen dioxide. It 
is an unstable yellow liquid, which slowly decomposes into hydrogen 
sulphide and sulphur. 

Hydrogen Selenide, H 2 Se, and Hydrogen Telluride, H 2 Te, are gases 
which closely resemble in chemical properties their analogue, hydrogen 
sulphide. 

Sulphur Monochloride, S 2 C1 2 , is made by passing chlorine 
over sulphur gently heated in a retort. The chloride distils 
over, and. is purified by redistillation. It is a yellow liquid, 
boiling at 138°. Its gas density has been found to be 68.8, 
showing that its molecule is S 2 C1 2 . It dissolves two thirds of 
its weight of sulphur. Sulphur monochloride gradually de- 
composes in contact with water, thus: 

2S 2 C1 2 + 2H 2 = 4HC1 + S0 2 -h 3S. 
Sulphur chloride is employed m vulcanizing rubber. 

Sulphur Dichloride, SC1 2 , is formed when chlorine acts on 
the monochloride at a temperature of 6° to 10°. It decom- 
poses on heating into monochloride and chlorine, 



152 THE SIXTH GROUP. 

Sulphur Tetrachloride, SC1 4 , is formed when chlorine is 
passed into the dichloride cooled to — 22°. The liquid ob- 
tained evolves chlorine as soon as the temperature rises. 

The compounds S 2 Br 2 , S 2 I 2 , and SI 6 have been obtained. 

A number of halides of selenium and tellurium are known. 
Tellurium tetrachloride boils without decomposition at 380°. 
The density of its gas has been found to be 130 at 448°; 
theory requires 133 for TeCl 4 . The tetrachlorides of sulphur 
and selenium dissociate on heating. 



The Sulphides and Hydrosulphides of the First 

Group. 

Potassium Hydrosulphide, K-S-H. — A solution of this 
substance is obtained by saturating a solution of potassium 
hydroxide with hydrogen sulphide. The liquid is colorless, 
strongly alkaline, and on evaporation out of contact with air 
deposits colorless crystals of the compound 2KSH + H,0. 
A solution of potassium hydrosulphide decomposes gradually 
on exposure to air, becoming yellow, and after a time colorless. 

Potassium Monosulphide, K 2 S or K-S-K, is obtained by 
saturating one half of a solution of potassium hydroxide with 
hydrogen sulphide and then adding the remaining half of the 
solution of potassium hydroxide. The solution absorbs oxygen 
from air, and becomes yellow owing to formation of higher 
sulphides. 

Potassium trisulphide, K 2 S 3 ; potassium tetrasulphide, K 2 S 4 ; 
and potassium pentasulphide. K 2 S 5 , are known. The latter 
compound is easily obtained by heating a solution of any of 
the other sulphides with flowers of sulphur. 



SULPHIDES OF COPPER AND SILVER. 153 

The potassium sulphides are decomposed by acids with 
evolution of hydrogen sulphide, and excepting KSH and 
K 2 S, with separation of sulphur, as shown in Exp. 103. 

Sodium Sulphides are similar to the potassium sulphides and 
are prepared by the same methods. 

Cu 
Cuprous Sulphide, Cu 2 S or i >S, is formed when copper 

Cu 

is burned in sulphur, and remains as a brittle dark gray mass 

when cupric sulphide is strongly ignited in hydrogen. Native 

cuprous sulphide or chalcocite occurs in handsome crystals of 

a blackish lead-gray color. It is an important ore of copper. 

Exp. 110. — Place in a narrow glass tube six inches long and closed at 
one end some sulphur and a copper wire. Heat strongly until the copper 
ceases to glow in the sulphur vapor. When cool examine the product. 

Cupric Sulphide, Cu — S, is obtained by the action of hydro- 
gen sulphide on cold solutions of cupric salts as a bulky black 
precipitate (see Exp. 108, a). While moist it readily oxidizes, 
with formation of cupric sulphate. If the precipitation is made 
in a boiling solution the cupric sulphide is denser, and easier 
to filter. The precipitated sulphide becomes blue on gentle 
ignition out of contact with air, and at higher temperatures 
loses half its sulphur. 

Silver Sulphide, Ag 2 S or Ag-S-Ag. — The native sulphide, 
known as argentite, occurs widely distributed as an ore. 
Silver sulphide is formed by the direct union of its elements, 
and by the action of hydrogen sulphide upon the metal. It 
separates as a blackish-brown precipitate on addition of 
hydrogen sulphide to a solution of a silver salt. 

Exp. 111. — Expose a clean silver coin to hydrogen sulphide gas or 
solution. The reaction is represented by the equation — 
2Ag + II-S-1I = Ag-S-Ag + H* 



154 THE SIXTH GKOUP. 

Gold Sulphides. — When an acid or neutral solution of gold 
is treated with hydrogen sulphide a black precipitate results, 
which, according to some experimenters, is a mixture of gold 
sulphide or sulphides with free sulphur and occasionally 
metallic gold. The precipitate is soluble in solutions of alkali 
sulphides. When sodium sulphide is used the following com- 
pound is formed : 

Sodium Aurous Sulphide, Na-Au = S + 4H 2 0, is prepared 
by strongly heating gold with sodium sulphide, treating the 
fused mass with water, and evaporating the solution out of con- 
tact with oxygen. It forms colorless monoclinic prisms, which 
become brown in the air. 



The Oxides and Hydroxides of Sulphur^ Selenium^ 
and Tellurium. 

Four compounds of sulphur and oxygen are known, namely : 
the sesquioxide, S 2 3 ; the dioxide, S0 2 ; the trioxide, S0 3 ; 
and the heptoxide, S 2 7 . 

Sulphur Sesquioxide, So0 3 . — When flowers of sulphur are gradually 
added to sulphur trioxide, dark-blue drops of the sesquioxide form, 
which at once solidify. The sesquioxide decomposes slowly at ordinary 
temperature, more readily on warming, into sulphur dioxide and sul- 
phur. It dissolves in fuming sulphuric acid to a blue solution. 

Sulphur Heptoxide, S 2 7 , is formed when electric sparks are passed 
through a mixture of sulphur dioxide and oxygen. It separates in oily 
drops, which solidify at 0°. It decomposes spontaneously, quickly on 
warming, into sulphur trioxide and oxygen. 

Sulphur Dioxide, Sulphurous Oxide or Anhydride, S0 2 . — This 
substance is formed when sulphur burns in air or oxygen. 
That sulphur dioxide contains its own volume of oxygen 
may be shown by the following experiment : 



SULPHUR DIOXIDE. 



155 



Exp. 112. — Through the stopper of the apparatus, Fig. 75, pass two 
stout wires, one of which is fastened to the handle of the platinum spoon 
B, and the other is connected by means of a piece 
of blowpipe platinum wire with the inside of the 
bowl of the spoon, so that the small platinum wire 
may be made to glow by means of a current from 
a battery. A bit of sulphur is melted about the 
connecting wire in the spoon. The apparatus is 
tilled with dry oxygen by passing the gas through 
it for a few minutes. The stopper is then put in 
place and mercury poured in, the stopper being 
loosened to allow the mercury to stand at the same 
height in both limbs. The apparatus should 
stand some time after filling, so that it may have 
the temperature of the room. "When ready to try 
the experiment, add or draw off mercury till it 
stands at the same height in both limbs ; then draw 
off a third or more of the mercury into an empty 
beaker to rarefy the gas in the bulb, and close the 
opening G with a cork. Connect for an instant the wires in A with a 
battery. The glowing platinum wire will inflame the sulphur. After 
the flame has disappeared remove the cork from C and pour back the 
mercury which was drawn off to diminish pressure. When the appa- 
ratus has cooled to the temperature of the room it will be found that the 
gas occupies a little less space than the original oxygen. The slight 
diminution in volume is due to the fact that when sulphur burns in 
oxygen or air a small quantity of sulphur trioxide forms, which is not 
gaseous at ordinary temperature. The cloud seen in the bulb is not 
sulphur dioxide, but trioxide, or if moisture is present it is sulphuric 
acid. 




Fig. 



Sulphur dioxide is prepared in several ways. A good 
method for obtaining the pure gas depends upon the decom- 
position of hot sulphuric acid by copper, represented by the 
following equation: 



on 



o 



2S0 2 < + On = S0 2 + S0 2 <^>Cu + -11,0. 



on 



o 



A little copper sulphide is formed at the same time. In ease 



156 THE SIXTH GROUP. 

the presence of carbon dioxide is not objectionable, charcoal 
may be used to advantage: 

C + 2H 2 S0 4 = 2S0 2 + C0 2 + 2H 2 0. 

Sulphur dioxide is a colorless gas, which is easily liquefied 
when exposed to low temperature or pressure. The liquid 
boils at — 8°, and when a strong blast of air is directed against 
its surface a portion of the dioxide freezes to a snow-like mass, 
melting at about — 80°. The vapor of pure sulphur dioxide 
mixed with air is rather less irritating when inhaled than the 
fumes from burning sulphur, which are chiefly sulphur diox- 
ide, but also contain a little sulphuric acid formed at the same 
time. The observed gas density of sulphur dioxide is 32. 
Hence its molecular weight is 61. Exp. 112 shows that one 
volume of sulphur dioxide gas contains one volume of oxygen. 

One volume of sulphur dioxide gas weighs 32 

" " " oxygen weighs 16 

Weight of sulphur in one volume of sulphur dioxide. 16 

Hence sulphur dioxide is composed of equal weights of sul- 
phur and oxygen, and in the molecular weight, 61, there are 
32 weights of sulphur, the weight of one atom, and 32 weights 
of oxygen, the weight of two atoms. Therefore the molecular 
formula is S0 2 . 

Sulphur dioxide is used in bleaching silk, wool, and straw, 
in place of chlorine, which injures these materials. It is a 
valuable disinfectant, and for this purpose and also for bleach- 
ing it is obtained by burning sulphur. 

Exp. 113. — Heat a flask containing copper clippings and strong sul- 
phuric acid, and pass the gas through water to free it from sulphuric 
acid fume, and then into distilled water until a saturated solution is oh- 



SULPHUROUS ACID. 



157 



tained. The sulphur dioxide gas may be collected over mercury or by 
displacing air. If liquid dioxide is desired, the 
gas should be dried by passing it through con- 
centrated sulphuric acid, and then into a test- 
tube immersed in a freezing mixture of ice and 
salt. The apparatus shown in Fig. 76 is con- 
venient. The liquid dioxide which collects in A 
may be drawn off by opening the pinchcock B. 
The tube C is held in a freezing mixture in a 
beaker. It is filled about two thirds with liquid 
dioxide, and without removing it from the freez- 
ing mixture is sealed by drawing off the neck in 
the flame of a blast-lamp. A number of tubes 
may be successively filled and kept until needed. 
A tube containing the dioxide should be cooled 
before breaking open the narrow end. 

Exp. 114. — Water may be frozen in a red-hot 
dish as follows. Heat a platinum capsule two 
inches in diameter by means of a good lamp to 
bright redness, and drop into it cautiously from 
a small test-tube a little liquid sulphur dioxide ; 
then pour in about the same bulk of water. The 
capsule is at once seized with tongs and the ice 
thrown upon a piece of wood. The liquid diox- 
ide assumes the spheroidal state in the hot dish 
and does not boil, being separated from the metal 
by a layer of gaseous dioxide. When the water is added the agitation 
causes the liquid dioxide to boil violently, and the heat required for 
vaporization is taken in part from the water which is frozen. 

Exp. 115. — Place a little liquid sulphur dioxide on a large watch- 
glass, and direct a current of air from a bellows on it. The white, 
snow-like mass obtained will gradually disappear without melting. 




Fig. 76. 



Sulphurous Acid, H,S0 3 , S0(0H) 2 or H-SO. -OH. —Water 

dissolves about 40 volumes of sulphur dioxide at ordinary 
temperature. The acid has not been isolated, but is assumed 
to exist, because the solution reacts acid, and yields witli bases 
sulphites. Dry sulphur dioxide does not redden blue litmus 
paper. A solution of sulphurous acid gradually absorbs 
oxygen from the air, with formation of sulphuric acid, A 



158 THE SIXTH GROUP. 

number of reactions in which the same change is effected are 
given under Sulphuric Acid. 

The constitution of sulphurous acid has been the s abject 
of careful investigations, which cannot be given here. The 
results thus far obtained iDdicate that H-S0 2 -OH is the 
structural formula. According to this view, two acid sodium 
(or potassium) salts are possible, Na-S0 2 -OH and H-S0 2 
— ONa. Since it is not known which of these formulas repre- 
sent the constitution of hydrogen sodium sulphite, the em- 
pirical formula HNaS0 3 will be used. 

Hydrogen Sodium Sulphite, HNaS0 3 . — This salt separates 
when a cold concentrated solution of sodium carbonate is sat- 
urated with sulphur dioxide, carbon dioxide being evolved. 

Sodium Sulphite, Na 2 S0 3 . — To prepare this substance 
sodium carbonate is saturated with sulphur dioxide, and to 
the solution as much more sodium carbonate is added. The 
anhydrous salt separates when the saturated solution is heated. 

The potassium salts HKS0 3 and K 2 S0 3 are analogous to the 
sodium sulphites. 

The sulphites are decomposed by acids, with evolution of 
sulphur dioxide. 

Sulphur Trioxide or Sulphuric Anhydride, S0 3 or S0 2 =0 — 

When dry sulphur dioxide and oxygen are passed over hot 
platinum sponge the two gases unite to form sulphur trioxide. 
It is more easily obtained by gently warming fuming sulphuric 
acid in a retort, and receiving the vapors in a flask. Thus 
prepared it contains a trace of water, melts at. 16°, and boils 
at about 46°. On standing at ordinary temperature long crys- 
talline fibres slowly form, which fuse above 50°, and change 
to the first-mentioned form. According to Weber, perfectly 
pure sulphur trioxide fuses at 14°. 8 and boils at 46°.2, and 
is changed by a small amount of moisture into the fibrous 
modification. 



SULPHUEIC ACID. 159 

Sulphur trioxide fumes in the air, and unites with water to 
form sulphuric acid. Its gas density corresponds to the 
molecule S0 3 . 

OH 

Sulphuric Acid, H 2 S0 4 or S0 2 <q H . — This acid is more ex- 
tensively used than any other acid. It is manufactured on 
an enormous scale by the so-called lead-chamber process. 
Sulphur dioxide, from burning sulphur or iron pyrites, FeS 2 , 
is passed into large lead chambers, into which are also con- 
ducted air, steam, and fumes of nitric acid. The complicated 
changes which occur may in part be represented as follows: 

Sulphur dioxide is converted into sulphuric acid by the nitric 
acid, which is reduced to nitrogen dioxide, N0 2 , thus: 

SO„ + 2NO a -OH = S0 2 <°| + 2N0 2 . 

The nitrogen dioxide in presence of water converts another 
portion of sulphur dioxide into sulphuric acid : 

S0 2 + N0 2 + H 3 = SO a <°g + NO. 



The nitrogen monoxide takes oxygen from the air supplied, 
and is changed to the dioxide : 

NO+O = N0 2 . 

Thus small quantities of the oxides of nitrogen serve as car- 
riers of oxygen, and suffice for the conversion of a largo quan- 
tity of sulphur dioxide into sulphuric acid. The dilute 
chamber acid is evaporated in lead pans until it contains about 
78 per cent of H a S0 4 , stronger acid attacking the lead. It 
is then further concentrated in glass or platinum retorts. The 
crude concentrated acid is commonly called oil of vitriol. 



160 THE SIXTH GROUP. 

Sulphuric acid may be detected by means of barium 

CI 
chloride, Ba<™, which forms with it and also with solu- 
tions of sulphates a white precipitate of barium sulphate, 
S0 2 <r\> Ba, which is insoluble in acids. 



Exp. 116. — Dilute a few drops of sulphuric acid largely with water 
and add a solution of barium chloride. Try to dissolve the precipitate 
obtained with hydrochloric acid. Formulate the reaction. 

Exp. 117. — Burn sulphur in a jar of air, or better, oxygen, and then 
pour in some barium chloride. The slight turbidity indicates that a 
little sulphuric acid is formed when sulphur burns. Next place in the 
jar a strip of wood moistened with concentrated nitric acid. Red fumes 
will be seen, also more barium sulphate. 

Exp. 118. — To sulphur dioxide obtained by burning sulphur in oxygen 
add bromine water. The color of the bromine will disappear. 






SO. 



OH 



+ Br 2 + 2H 2 = S0 2 <™ + 2HBr. 



Test for the presence of sulphuric acid. Chlorine in presence of water 
also converts sulphur dioxide into sulphuric acid. 

The oil of vitriol of commerce is frequently dark-colored 
from organic matter, contains lead sulphate, and often other 
impurities. It is purified by distilling in glass. The most 
concentrated acid that can be obtained by distillation contains 
about 2 per cent of water. If such acid is cooled, crystals of 
H. 2 S0 4 separate, which are free from water. These melt at 
10°. 5, and remain liquid when cooled below this temperature. 
The pure liquid acid has a density of 1.854 at 0°. Ordinary 
concentrated sulphuric acid or oil of vitriol having a density of 
1.819 at 15° contains 89. 7 per cent of acid. The pure acid 
fumes when warmed, part of the H 2 S0 4 being dissociated into 
sulphur trioxide and water, and on heating to 338°, the boil- 
ing point of the liquid, an acid remains containing 1.2 to 1.6 
per cent of water, and which distils unchanged. 



SULPHURIC ACID. 161 

The density of the vapor obtained by heating sulphuric acid 
to 440° has been found to be 25, which corresponds to that of 
a mixture of equal molecules of water and sulphur trioxide 
and shows that sulphuric acid does not exist in the gaseous 
state. 

When sulphuric acid is mixed with water much heat is 
evolved, the mixture at the same time undergoing contraction. 
The acid absorbs moisture rapidly from the air, and is therefore 
used for drying gases. Oil of vitriol chars wood and paper, 
and burns the skin. 

Exp. 119. — Pour a thin stream of oil of vitriol into water. The heat- 
ing of the mixture will be evident. Water should not be poured into 
the acid, owing to the danger of the acid being thrown out by sudden 
bursts of steam. If oil of vitriol contains lead sulphate, this will sep- 
arate as a white precipitate when the acid is diluted with water. 

Exp. 120. — Place in a test-tube oil of vitriol and pieces of cloth, paper, 
and wood, and after some days observe the result. 



BASES, ACIDS, AND SALTS. 

The oxides and hydroxides of the metals possess, as a rule, 
the property of neutralizing acids. This peculiar property 
is termed basic. The oxides which exhibit it are called basic 
oxides, and the hydroxides are called bases. Later we shall 
learn that some metallic hydroxides are acids. Soluble bases 
turn red litmus blue, having what is called an alkaline reac- 
tion. Acids, on the contrary, redden blue litmus, and decom- 
pose carbonates with evolution of carbon dioxide. There are 
many compounds which are neither acids, bases, nor salts ; but 
the distinguishing characteristics of strong bases and acids are 
easily recognized. 

Let us next consider the constitution of two bases with 
which we are already familiar, namely, potassium hydroxide 
and sodium hydroxide. 

If we regard oxygen as bivalent and potassium and hydrogen 
as univalent, we may represent potassium hydroxide by the 
structural formula K-O-H. Sodium hydroxide may be 
likewise represented as Xa-O-H. The group OH is called 
hydroxyl, and has already been noticed under hydrogen 
dioxide. It is an example of a compound radical. 

Radicals. — A simple radical is a single atom. A compound 
radical is a group of atoms which deports itself like a single 
atom, and which maybe transferred from one compound to 
another without losing its identity. Thus OH may be regarded 
in K-OH as replacing 01 in K-Cl. Compound as well as 
simple radicals possess valence. Later we shall become ac- 
quainted with a number of compound radicals. 



BASES, ACIDS, AND SALTS. 163 

Chemists regard oxygen acids as hydroxides, that is, com- 
pounds containing the radical OH. The structural formula 

OH 

of sulphuric acid is^S0 2 <Qjr. This represents the acid as 

containing the radicals S0 2 and OH. Later it will be seen 
that the properties of sulphuric acid are best indicated by the 
structural formula, which is therefore more useful than the 
simpler one, H 2 S0 4 . 

In the electrolysis of sodium chloride sodium appears at 
the negative electrode and is therefore electro-positive, and 
chlorine is disengaged at the positive electrode and is electro- 
negative. A solution of potassium sulphate is decomposed by 
electricity, with formation of potassium hydroxide at the 
negative and sulphuric acid at the positive electrode. Since 
salts are thus separated into positive and negative parts in 
electrical relations, the former being basic and the latter 
acidic, it is customary to call basic radicals positive radicals, 
and acid radicals negative radicals. Potassium chloride, 
KC1, contains the positive radical K, the same that in KOH 
gives it basic properties. Hydrogen is common to both acids 
and bases, and is regarded as possessing neither acidic nor 
basic properties. Assuming that sulphuric acid contains two 
hydroxyls, the negative radical is either S, SO or S0 2 ; the 
latter is considered to be the acid radical, and is termed sul- 
phuryl. 

Salts are regarded as built on the same plan as the acids 
from which they are derived. Thus the structural formula of 

potassium sulphate is S0 2 <Q~jr-. Nitric acid is NO a -0-B 

and potassium nitrate is N0 2 -0-K, N0 2 being the acid 
radical. The formulas in the following equations of the 
formation of salts represent similarity in the structure of 
oxygen bases, acids, and salts, in this respect, that in these 
three classes of compounds radicals arc linked together by 
oxygen: 



164 BASES, ACIDS, SALTS. 

K-O-H + N0 3 -0-H = X0 2 -0-K + H-O-H; 
K -0-H + S0 2 <^:g = S0 2 <^:§ + H-0-H; 



2K-0-H + SO a <^ = S0 2 <°_| + 2H-0-H. 



From the structural formulas of bases, acids, aud salts the 
following definitions are derived : 

An oxygen base consists of a positive (basic) radical joined 
by oxygen to hydrogen. 

An oxygen acid consists of a negative (acid) radical joined 
by oxygen to hydrogen. 

An oxygen salt consists of a positive radical or radicals 
joined by oxygen to a negative radical or radicals. 

The formation of salts by the reaction between acids and 
bases is always accompanied by the elimination of water. 

Sulphuric acid contains two atoms of hydrogen, which may 
be replaced by metals. This property is indicated by calling 
it a dibasic acid. Two molecules of potassium hydroxide, 
K-O-H, are required to saturate it, and but one is needed 
to saturate monobasic nitric acid, N0 2 -0-H. In the same 
way potassium hydroxide is considered a monacid base; cal- 
cium hydroxide (slaked lime) a diacid base, since it reacts 
with two molecules of a monobasic acid, or one of a dibasic 
acid, to form a salt, thus: 



Ca< 0-H + 2H " C1 = Ca< Cl +2H-0-H; 

Ca<£:|-f 2M) 2 -0-H = Ca<^:|^ + 2H-0-H; 

Ca< 0-H+ S0 »<0-H = Ca<^>S0 2 + 2H-0-H. 



BASES, ACIDS, AKD SALTS. 165 

It is evident that the basicity of an acid is indicated by the 
valence of the acid radical, and the acidity of a base by the 
valence of the basic radical. The same may be expressed dif- 
ferently, as follows: 

An acid is mono-, di-, tri-, or more basic according as it 
contains one, two, three, or more atoms of hydrogen joined by 
oxygen to the acid radical. 

A base is mono-, di-, tri-, or more acid according as it 
contains one, two, three, or more hydrogen atoms joined by 
oxygen to the basic radical. 

Normal Salts are those in which the bonds of the acid 
radical or radicals equal the bonds of the basic radical or 
radicals. Potassium nitrate, N0 2 -0-K, and potassium sul- 
phate, S0 2 <q_t^-, are examples. 

Acid Salts are those in which the bonds of the acid radical 
exceed the bonds of the basic radical or radicals, as for ex- 

O— TC 

ample in hydrogen potassium sulphate, S0 2 < n tt- 



Basic Salts are those in which the bonds of the basic radical 
or radicals exceed the bonds of the acid radical. For instance, 

basic lead nitrate, t| _Q>Pb. 



Acid Anhydrides. — Oxides which combine with water to 
form acids are termed acid anhydrides. For example, sul- 
phuric anhydride, S0 3 , and water unite to form sulphuric 
acid. The acid anhydrides are derived from acids by the re- 
moval of the elements of water: 



II a S0 4 - H a O = SO, 



166 the' sixth group. 

Salts result from a number of reactions of a different char- 
acter from those already stated. They are formed by the 
replacement of the hydrogen of an acid by a metal, 

Zn + SO,<g:| = S0 2 <°>Zn + H a ; 

by the combination of an acid anhydride with a metallic oxide, 

S0 2 = + Ca=0 == S0 2 < Q>Ca; 
by the reaction between a metallic oxide and an acid, 

Ca=0 + S0 2 < O-H = S0 2 <°> Ca + H-O-H; 
and also by an exchange of radicals between two salts, 

SO » < 0-K + Ba< Cl = s ° 2 <o> Ba + 2K - CL 



Sulphuryl Hydroxychloride or Sulphuric Hydrochloride, 

CI 
S0 2 <qjj. — This compound is formed by the direct union of 

sulphuric anhydride and hydrochloric acid. It is best ob- 
tained by heating a mixture of sulphuric acid and phosphorus 
oxy chloride : 

2SO *<OH + P0C1 3 = 2S °.<OH + HP0 * + HCL 

Phosphorus oxychloride is used to effect the replacement of 
hydroxyl by chlorine in a number of reactions. Sulphuryl 
hydroxychloride is a colorless fuming liquid, boiling at 158°. 



CONSTITUTION OF SULPHURIC ACID. 167 

Its gas density accords with its formula. It is decomposed 
violently by water, with formation of sulphuric and hydro- 
chloric acids. Hence it may be viewed as a chloranhydride 
of sulphuric acid. 

CI 
Sulphonyl Chloride or Sulphuric Dichloride, S0 2 < C1< — Sul- 
phur dioxide and chlorine unite in sunlight to form this 
substance. It is, however, best prepared by heating sulphuryl 
hydroxychloride in a sealed tube to 180° for twelve hours: 

2S0 1 <^ H =SO,<gJ + SO i <g. 

Sulphuryl chloride is a colorless liquid, boils at 70°, and has 
a gas density of 67.5. It yields with little water sulphuryl 
hydroxychloride and hydrochloric acid, and with an excess of 
water sulphuric and hydrochloric acids. It is, therefore, a 
chloranhydride of sulphuric acid. 

Constitution of Sulphuric Acid. — The chloranhydrides of 
sulphuric acid are formed, as already shown, by the replace- 
ment of oxygen and hydrogen by chlorine. The simplest 
supposition is that the acid contains the radical hydroxyl, 
OH; a view supported by the reactions of the chloranhydrides 
with water, whereby an atom of chlorine is replaced by the 
elements of hydroxyl. The changes may be represented thus: 



so 2 <g±Jil- - H = so 2 <o r H + H-ci. 

SO><PTH|-0-H = S0 2 <0 : H + H-01. 



Both changes may occur simultaneously: 
S0 2 < 



01 , II 
Ol+ll 



l_ 



£u = s°.<cmI + •'"-«• 



168 



THE SIXTH GROUP. 



It may be stated that there are many reactions analagous to 
the foregoing in which chlorine replaces hydroxyl, and that 
proof of the presence of hydroxyl in a compound is found in 
such reactions. 

If we accept the structural formula of sulphuryl chloride we 
see that the foregoing equations are simple expressions of the 

OH 

nature of the reactions. Further, the formula S0 2 <qtt 

represents the relation of sulphuric acid to the chloranhy- 
drides derived from it. The mutual relation between the 
sulphur atom and two of the oxygen atoms is different from 
that of the oxygen atoms in the hydroxyl group, and these 
last two are influenced by or united to both sulphur and 
hydrogen. These supposed differences in the relations of 
the atoms in the molecule are indicated by the formula 
O^/O-H 

o^<o-ir 



Sulphur Hexhydroxide or Ortho sulphuric Acid, S(0H) 6 . — 
When sulphuric acid, H 2 S0 4 , is mixed with water in the pro- 
portion of one molecule of the former to two of the latter 
the contraction in bulk is greater than with other propor- 
tions. It is supposed that the compound H 6 S0 6 (= H 2 S0 4 
+ '3H 2 0) is formed. Its only structural formula appears to be 
OH 



A few salts are known which may be re- 



garded as derivatives of this acid; for example, the mercury 
0. 




>H * 



sulphate, S 




THIOSULPHUEIC ACID. 169 

Monometasulphuric Acid, S0(0H) 4 . — When a mixture of 
equal molecules of sulphuric acid and water is cooled, crystals 
having the composition H 2 S0 4 + H a O form. These melt at 7°. 5. 
Monometasulphuric acid is derived from orthosulphuric acid 
by the removal of one molecule of water, giving the formula 

OH 

OH 

OH . By taking away another molecule of 



water, ordinary or dimetasulphuric acid results, and this in 
turn yields sulphuric anhydride by further dehydration. 
This hypothesis of the derivation of acid hydroxides accords 
with the theory of valence, and is a help in studying them. 

SO <0H 

Disulphuric Acid, HS 2 7 or 2 . — This acid is formed 

S0 2 <0H 

from sulphuric acid by the removal of water, thus: 

S ° 8< OH _ SO,<OH 

SO *<OH ~ S ° a< H ' 

Disulphuric acid forms well-defined salts, is dibasic, fumes 
in the air, and decomposes on warming into sulphur trioxide 
and sulphuric acid. The fuming sulphuric acid of commerce 
is a mixture of sulphuric and disulphuric acids. It is manu- 
factured from iron pyrites, FeS a , by a process of roasting 
and distilling, and is also obtained by adding sulphur trioxide 
to sulphuric acid. The chloranhydride of disulphuric acid, 

so/ 01 

SQ 2> ^ » * s known. 



■\01 

Chiosi 

name indicates, may be regarded as sulphuric acid with an 



ATT 

Thiosulphuric Acid, H a S a 3 or S0 9 <gg.— This acid, as its 



170 THE SIXTH GKOTTP. 

oxygen atom replaced by sulphur. The free acid is unknown, 
as it immediately decomposes when liberated from its salts 
into sulphur and sulphur dioxide, thus : 

S0 2 <°g = H 2 + S0 a + S. 

Sodium Thiosulphate, Na 2 S 2 3 + 5H 2 or BO,<2jf* + 

5H„0. — This salt is more commonly known as hyposulphite 
of soda. It is obtained by boiling a solution of sodium sul- 
phite with flowers of sulphur, when the following reaction 
occurs : 

Na n SO. + S = KaJSLO.. 



2~2 



Other methods are also used in manufacturing it. 

Sodium thiosulphate forms large transparent monoclinic 
crystals, which are permanent in the air and very soluble in 
water. It is used in the manufacture of paper as an antichlor, 
in photography, and as a reagent in the laboratory. 

TT 

Hyposulphurous Acid, H 2 S0 2 or SO< QH .— Zinc dissolves in a solution 

of sulphurous acid without evolution of lrydrogen, and with formation 
of hyposulphurous acid. The yellow solution which results possesses 
greater bleaching power than sulphurous acid. If a solution of acid 
sodium sulphite is treated with zinc, there is formed, together with 

TT 

other products, sodium hyposulphite, SO<Q-^r . This compound is 

used in calico-printing. From the hyposulphite the free acid is liber- 
ated by oxalic acid. The solution obtained has a deep orange color, 
and decomposes quickly with separation of sulphur. The acid is mono- 
basic. Solutions of its salts on exposure to air absorb oxygen with 
formation of sulphites. 

DitMonic Acid, H 2 S 2 O e ; Trithionic Acid, H 2 S 3 6 ; Tetrathionic Acid, 
H 2 S 4 6 ; and PentatMonic Acid, H 2 S 5 6 , have long been known. 

Selenium Dioxide, Se0 2 . — This is the only known oxide of 
selenium. It is prepared by burning selenium in a current of 



SELENIOUS ACID — TELLUROUS ACID. 171 

dry oxygen. It forms needle-like crystals, which volatilize 
without fusion at about 300°. The dioxide is the acid anhy- 
dride of selenious acid. 

OH 

Selenious Acid, H 2 Se0 3 or SeO<Q-rr, separates in colorless 

crystals on cooling a solution of five parts of selenium dioxide 
in one part of hot water. It forms salts which are analogous 
to the sulphites. It also forms compounds with an acid 
selenite, such as, HKSe0 3 -f H 2 S0 3 . 

OH 

Selenic Acid, H 2 Se0 4 or Se0 2 <QTT, is prepared by decom- 
posing silver selenite, suspended in water, by bromine : 

Ag,SeO, + Br, + H 3 = H s SeO, + 2AgBr. 

The insoluble silver bromide is removed, and the solution is 
concentrated by evaporation. 

The acid dissolves zinc and iron with evolution of hydrogen, 
and when hgt and concentrated it dissolves gold with reduc- 
tion to selenious acid. 

Tellurium Dioxide, Te0 2 , is formed when tellurium burns in 
air. It is but slightly soluble in water, and the solution does 
not react acid to litmus. It possesses feeble basic as we 1 ! as 
acidic properties. 

Tellurium Trioxide, Te0 3 , results when telluric acid is 
heated to a temperature below a red heat. It has an orange- 
yellow color, and is insoluble in water, in cold hydrochloric 
acid, and a solution of potassium hydroxide, unless the latter 
is hot and concentrated. 

OH 
Tellurous Acid, H a TeO„ or Te0< o -rr, separates when a 

solution of tellurium chloride is poured into water, in which 



172 SULPHATES OF THE FIRST GROUP. 

the acid is but slightly soluble. The salts K 2 Te0 3 , HKTe0 3 , 
and K 2 Te0 3 .3Te0 3 , are examples of tellurites. 

Telluric Acid. — When a dilute solution of this acid is 
evaporated, crystals having the composition of Te(OH) 6 sepa- 
rate. This is the analogue of orthosulphuric acid. The 
crystals lose water at 160°, and change into H 2 Te0 4 . Telluric 
acid forms salts analogous to the sulphates, and also such acid 
compounds as K 2 Te0 4 +3Te0 3 + 4E.fi. 



Sulphates of the First Group. 

Sodium Sulphate, Na 2 S0 4 or S0 s <q_JJ* — This salt is made 

commercially in enormous quantities by heating common salt 
with sulphuric acid. It is also obtained in the manufacture 
of nitric acid from sodium nitrate. It occurs in many natural 
waters. The crude sodium sulphate of commerce is known 
as " salt cake," and is one of the intermediate products in the 
manufacture of sodium carbonate from sodium chloride. It 
is used in glass-making. 

When a solution of sodium sulphate evaporates at ordinary 
temperature, the decahydrated salt Na 2 S0 4 -J- 10H 2 O, com- 
monly known as Glauber's salt, separates in colorless mono- 
clinic crystals. The crystals effloresce in dry air, melt at 33°, 
and give up all their water below 100°. The heptahydrated 
salt ]^a 2 S0 4 -f- TH 2 sometimes separates in hard rhombic crys- 
tals when a supersaturated solution of sodium sulphate is 
cooled. This hydrate is sometimes sold for Glauber's salt. 

Sodium sulphate dissolves to the greatest extent in water 
when the mixture is warmed to 34°. at which temperature a 
saturated solution contains 50 parts of Xa 2 S0 4 in 100 parts of 



POTASSIUM SULPHATE. 173 

water. On heating such, a solution monohydrated sodium 
sulphate will separate. Pulverized Glauber's salt and strong 
hydrochloric acid make a convenient freezing mixture. The 
salt Na 2 SO, -f- ?H 2 will not answer instead of Glauber's salt. 



Exp. 121.— Make a solution by heating to 84° on a water-bath a mix- 
ture of,125cc. of water and \ k. of Glauber's salt in a flask. Shake 
frequently, and when no more appears to dissolve filter the solution into a 
flask. Cover the neck of the flask with paper to keep out dust. Allow 
the solution to cool to common temperature, when it will be supersaturated, 
that is, will contain more of the salt than is ordinarily soluble. Next drop 
into the solution a fragment of Glauber's salt. Crystals will grow very 
rapidly. Dust from the air will also cause crystals to form, owing, per- 
haps, to minute particles of sodium sulphate in the air. 



Hydrogen Sodium Sulphate, Acid Sodium Sulphate, HKaS0 4 

O-Na 
or S0 2 <q_tt . — This salt separates in crystals when a solu- 
tion of equal molecules of sulphuric acid and sodium sulphate 
is evaporated by heat. 

so /°- Na 

Sodium Disulphate or Pyrosulphate, Na o S,0 7 or cn ">0 

bU *\0-Na 
— This compound results from heating the preceding salt. It 
is best prepared by heating a mixture of sodium sulphate and 
sulphuric acid until the mass fuses quietly at a dull-red heat. 
At higher temperatures S0 3 is evolved, leaving Na 3 S0 4 . 



0— K 

Potassium Sulphate, K 2 S0 4 or S0 2 < _- K -, crystallizes in 

small rhombic pyramids. The salt is soluble in 10 parts of 
water at ordinary temperature. It is obtained from certain 
minerals and as a by-product in some chemical manufactures. 
It is used in making potash-alum, and other potassium com- 
pounds, 



174 SULPHATES OF THE FIRST GKOUP. 

0— K 

Hydrogen Potassium Sulphate, HKS0 4 or S0 2 < Q „, maybe 

prepared by melting together 13 parts of potassium sulphate 
with 8 parts of oil of vitriol. It is soluble in about half its 
weight of boiling water, and the solution on cooling deposits 
rhombic crystals. It is decomposed by a large quantity of 
water, and the dilute solution when evaporated deposits first 
crystals of the normal salt, then crystals having the compo- 
sition K 2 S0 4 + HKSO t , and finally hydrogen potassium sul- 
phate. 

Potassium Disulphate, K 2 S 2 7 , is similar in properties to 
sodium disulphate, and is prepared in a like manner. 

so /°- K 

Hydrogen Potassium Disulphate, HKS 2 7 or Qn 2 >0 . — 

&U2 \OH 
This compound is obtained in crystals by dissolving potassium 
disulphate in fuming sulphuric acid. 

Hydrogen potassium sulphate and potassium disulphate and 
the corresponding sodium salts are valuable as fluxes. At a 
red heat they convert into sulphates oxides or salts which are 
not affected by boiling sulphuric acid. 



Lithium Sulphate crystallizes with one molecule of water, while 
Rubidium aud Caesium Sulphates, like the potassium sulphate, form only 
anhydrous salts. The salts, HLiS0 4 , HRbSQ 4 , and HCsS0 4 are known. 

Cupric Sulphate, Copper Sulphate, CuS0 4 or S0 2 < ft >Cu. 

— This salt is manufactured on a large scale by roasting ores 
containing copper sulphide and lixiviating with water. The 
solution yields 011 evaporation blue crystals having the compo- 
sition CuS0 4 + 5H s O, which are known in commerce as blue 
vitriol and blue-stone. It is also made by dissolving copper 



SILVER SULPHATE. 175 

scale in sulphuric acid. The crystallized salt loses four mole- 
cules of water at 100°, becoming greenish white, and the last 
molecule at 200°, becoming pure white in color. The com- 
position of the monohydrate is that of inonometa-sulphate, 

Tr r) >S0< o >Cu. The anhydrous CuS0 4 absorbs water on 

exposure to moist air, and becomes blue. When intensely 
ignited it is changed into cupric oxide. Both the hydrous 
and anhydrous salt absorb hydrochloric acid gas with forma- 
tion of cupric chloride and sulphuric acid. This reaction is 
made use of in separating hydrochloric acid gas from other 
gases. One part of CuS0 4 -j- 5H 2 is soluble in about 2-J 
parts of water at ordinary temperature, and one half a part at 
100°. The salt is insoluble in alcohol. Blue vitriol is the most 
common of copper salts. It is used in dyeing, in preparing 
other copper compounds, in electroplating, and electrotyping, 
and in galvanic batteries. 



Exp. 122. — Heat powdered blue vitriol until the residue is white. Ex- 
pose a portion of the anhydrous sulphate to the air, add another portion to 
common alcohol, and agitate. In both cases the sulphate will become 
blue— a proof of the presence of water in the air and alcohol. 



Silver Sulphate, Ag,S0 4 or S0 2 <q_^|.— Silver like copper 

dissolves in hot oil of vitriol with the evolution of sulphur di- 
oxide. This fact is applied in separating gold from silver in 
alloys, the silver dissolving in the sulphuric acid and the gold 
remaining. Silver sulphate requires 88 parts of boiling water 
and about 200 parts of cold water to dissolve it. It is much 
more soluble in water containing nitric acid or sulphuric acid. 
Silver sulphate separates as a white powder when sulphuric 
acid is added to a not too dilute solution of silver nitrate. It 
fuses at a dull-red heat, and at a higher temperature decom- 
poses into silver, oxygen, and sulphur dioxide. 



176 SULPHATES OF THE FIRST GROUP. 

Hydrogen Auryl Sulphate, HAuOS0 4 or 0=Au-0-S0 2 0H. 

— This compound is formed when auryl nitrate or hydrox- 
ide is digested at 200° with concentrated sulphuric acid. The 
crystalline meal which separates is freed from the mother- 
liquor on a porous plate, washed with nitric acid, density 1.49, 
and dried over lime. The salt has a canary-yellow color, is 
very hygroscopic, and is decomposed by water into auryl 
hydroxide and sulphuric acid. The hydrogen of this salt may 
be replaced by potassium forming the compound KAuOS0 4 . 

Gold Sulphate, AuS0 4 , is formed by the rapid evaporation of 
a solution of auryl sulphate at 250°, and separates in scarlet- 
red prisms which absorb moisture and become black. The 
salt is decomposed by water with formation of the compound 
Au 6 H 4 8 = 6AuO+2H 2 0. 



CHROMIUM, MOLYBDENUM, TUNGSTEN, 
AND URANIUM. 

Chromium^ Cr. 

Atomic Weight, 52.5. Density, 6.8. 

Chromium was discovered independently by Vauquelin and 
Klaproth in 1797. The chief source of chromium is chromite, 
FeCr 2 4 . Metallic chromium is prepared by intensely heating 
the oxide with charcoal;, or by reducing the chloride with zinc 
at high temperatures. It has been obtained by these methods 
only in the form of powder or small crystals. The metal is 
deposited in the electrolysis of its chloride as glistening scales 
or a coherent plate. It is more infusible than platinum, and as 
hard as corundum. The pure metal is more permanent in 
damp air than iron. It is soluble in hydrochloric and sulphuric 
acid, but not in nitric acid. 

Chromium forms compounds corresponding to CrO, Cr o 3 , 
and Cr0 3 . The first two are basic oxides, and the last is an 
acid anhydride. 



Chromous Compounds. 

CI 
Chromous Chloride, CrCl 2 or Cr< c ,. — This compound is ob- 
tained by heating chromic chloride in hydrogen, and also by dis- 
solving metallic chromium in hydrochloric acid. Chromous 
chloride oxidizes readily, and is an energetic reducing agent. 
12 



178 THE SIXTH GBOUP. 

Chrornous Hydroxide, Cr(0H) 2 , separates when potassium 
hydroxide is added to a solution of chrornous chloride. It 
absorbs oxygen on exposure to air. When heated it is con- 
verted into chromic oxide with liberation of hydrogen: 

2Cr(OH) 2 - Cr 2 3 + H 2 + H 2 0. 

Chrornous Sulphate, CrS0 4 + 7H 2 0.— This salt forms bhie 
crystals. Like other chrornous compounds it quickly oxidizes 
in the air. 



Chromic Compounds. 

Solutions of chromic salts have a violet color, which changes 
to green on heating. Green solutions become violet on long 
standing at common temperature. The different colors of 
the solutions have been ascribed to two modifications of 
chromic salts. Eecent investigations indicate that when a 
violet solution of a chromic salt becomes green on heating, 
the normal salt decomposes into a basic salt and free acid. 

Chromium Sesquioxide, Chromic Oxide, Cr 2 3 or = Cr-O- 

Cr = 0. — This oxide is obtained by igniting chromic hy- 
droxide as an amorphous green powder, and in dark -green 
crystals by passing chromyl chloride through a red-hot tube. 
Chromium sesquioxide fuses in the flame of the compound 
blowpipe, and solidifies on cooling to a dark crystalline mass, 
which is sufficiently hard to scratch glass and steel. 

Chromium sesquioxide is used to impart a green tint to 
glass and porcelain, and as a green pigment in paints. 

It dissolves slowly in hot sulphuric acid, and is insoluble in 
other acids. When ignited with alkalies in contact with air 
it absorbs oxygen with formation of chromates. 



CHEOMIC COMPOUNDS. 179 

Chromic Hydroxides. — Ammonia produces in a green solu- 
tion of a chromic salt a grayish-green and in a violet solu- 
tion a grayish-blue precipitate, which, after drying over sul- 
phuric acid, has the composition Cr(OH) 3 -f- 2H 2 0. This 
compound loses water readily, and is converted, as already 
stated, into anhydrous oxide by ignition. Potassium and 
sodium hydroxides throw down from a solution of a chromic 
salt a green hydroxide containing alkali. The precipitate 
dissolves in excess of alkali to a deep-green solution, from 
which on long standing or boiling chromic hydroxide sep- 
arates. 

Chromic Chloride, CrCl 3 , is obtained by passing chlorine 
over an intensely heated mixture of chromic oxide and char- 
coal. It is nearly insoluble in pure water, but dissolves easily 
to a green solution in water containing a little chromous 
chloride. A solution of chromic chloride is also obtained by 
dissolving chromic hydroxide in hydrochloric acid. 

so *<o)c r 

Chromic Sulphate, Cr 2 (S0 4 ), or S0 2 <^ .—Chromic hy- 

so 2 <«/ Cr 

droxide dissolves in strong sulphuric acid to a green solution, 
which does not yield a violet crystalline precipitate on adding 
alcohol. On long standing in moist air the solution absorbs 
water and changes to a crystalline mass, which dissolves in 
water to a violet solution, from which alcohol precipitates the 
hydrate Cr Q (SOJ, + 18H a O. Chromic sulphate and alkali 
sulphates form alums, of which the following is the most 
common. 

Potassium Chromic Sulphate, Chrome Alum, K a S0 4 .Cr(S0 4 ), 

-j- 24H 2 0. — This double salt is easily made by adding sul- 



180 THE SIXTH GROUP. 

phuric acid to a solution of potassium dichromate and some 
reducing agent, such as sulphur dioxide or alcohol. The 
solution on standing several days deposits ruby-red octahe- 
drons. The salt is used in dyeing. 



Compounds of Hexvalent Chromium. 

F \ / F 
Chromium Hexfluoride, CrF 6 or F-)Cr^-F. — This compound 

F/ \f 

is prepared by heating in a retort a mixture of lead chromate, 
fluor spar, and fuming sulphuric acid. The red vapor which 
distils over condenses to a red fuming liquid. 

/O 
Chromium Trioxide, CrO, or Cr=0. — This substance is 

Xo 

easiest obtained by treating a cold saturated solution of po- 
tassium dichromate with 1.5 volumes of concentrated sulphuric 
acid. The chromium trioxide separates on cooling in long 
dark-red crystals, which are sufficiently pure for many pur- 
poses. To prepare pure trioxide, chlorium hexfluoride is de- 
composed by water: 

CrF 6 -f4H 2 = Cr0 3 + H 2 + 6HF. 

The red solution leaves on evaporation an amorphous red mass 
of the oxide, Cr0 3 . 

Chromium trioxide melts at 193°, and at 250° is reduced to 
Cr 2 3 . It is one of the best oxidizing agents known. When 
alcohol is dropped upon it rapid combustion takes place. A 
solution of the trioxide in glacial acetic acid or sulphuric acid 
is often used to oxidize organic compounds, and a mixture of 



COMPOUNDS OF HEXVALENT CHKOMIUM. 181 

potassium dichromate and sulphuric acid is much used for 
cleansing glass and porcelain in the laboratory. 

OH 

Chromic Acid, H 2 Cr0 4 or Cr0 2 < OH . — This acid is supposed 

to have a constitution like that of sulphuric acid, because 
chromates are analogous to sulphates. It is obtained by the 
following process. Less water than is required to form H 2 Cr0 4 
is added to chromium trioxide, the mixture heated to 100°, 
the solution decanted, and cooled to 0°, when small red crys- 
tals of chromic acid separate. These crystals, after drying 
over oil of vitriol, have the composition H 2 Cr0 4 . The acid is 
very hygroscopic. 

OK 

Potassium Chromate, K 2 Cr0 4 or Cr0 2 < 0K -. — This salt is 

obtained by adding potassium hydroxide to potassium dichro- 
mate: 

>0 +2KOH = 2CrO a <x£ + fl a O. 

The solution on evaporation deposits at first red crystals of 
potassium dichromate, and afterwards yellow crystals of the 
chromate. Solutions of the chromate are decomposed by 
acids, even by carbonic acid, with formation of the dichro- 
mate. Potassium chromate is isomorphous with potassium 
sulphate. The salts Na 2 Cr0 4 + 10H 2 O and Na 2 S0 4 + 10H 9 O 
are also isomorphous. 

Lead Chromate, PbCr0 4 or Cr0 2 <Q>Pb — This salt sepa- 
rates as a yellow precipitate when solutions of a Lead salt and 
potassium dichromate are mixed. It is extensively used as a 
pigment known as chrome yellow. Lead chromate fuses at 
a red heat, and forms on cooling a compact brown mass, yield- 
ing a yellowish-brown powder. Above its melting point it 



182 THE SIXTH GROUP. 

loses a portion of its oxygen. It is a valuable reagent, much 
used in the combustion of carbon compounds. Yellow lead 
chromate is converted by caustic alkalies into a chrome red, 
having the composition Pb 2 Cr0 5 . 

Dichromates are analogous in constitution to the disul- 
CrO - OH 
phates. Dichromic acid, > , has not been obtained, 

CrO - OH 

CrO -OH 
but chrome sulphuric acid, > , is formed when Or0 3 

SO - OH 
is dissolved in H 2 S0 4 , just as disulphuric acid results from 
the union of S0 3 with H 2 S0 4 . 

CrO - OK 

Potassium Dichromate, K 2 Cr0 27 or >0 . — This impor- 

Cr0 2 - OK 

tant compound forms splendid red triclinic crystals, soluble 
in about ten parts of water at ordinary temperature, and more 
abundantly in boiling water. The salt fuses below a red heat, 
and at a higher temperature decomposes into oxygen, chro- 
mium sesquioxide, and potassium chromate. It is largely 
used in dyeing, and in the preparation of chromium com- 
pounds. 

A film of gelatin and potassium dichromate darkens on ex- 
posure to light, the former becoming insoluble, and the latter 
being reduced to chromium sesquioxide. These changes are 
made use of in a photographic printing process. 

Potassium dichromate is manufactured from chromic iron or chro- 
mite, FeCr 2 4 . A mixture of the ore, potassium carbonate, and lime 
is roasted; in the process oxygen is absorbed from the air, and potassium 
and calcium chromates are formed. The roasted mass is lixiviated with 
water, and potassium sulphate is added to the solution to precipitate the 
calcium as sulphate. To the hot concentrated solution sulphuric acid 
is added to convert the potassium chromate into dichromate, which 



COMPOUNDS OF HEXVALENT CHEOMIUM. 183 

separates in crystals as the solution cools. Other dichromates are 
(NH 4 ) 2 Cr 2 7 , Na a Cr a 7 +2H 2 0, Ag 2 Cr 2 7 , and BaCr 2 7 + 2H 2 0. 

The following polychromates have also been prepared: potassium 
trichromate, K 2 Cr 3 Oi ; tetrachromate, K 2 Cr 4 13 ; ammonia trichro- 
mate, CN"H 4 ) 2 Cr 3 Oio; and hexchromate, (NH 4 )2Cr 6 0i 9 . These com- 
pounds are viewed as similar in constitution to the dichromates, in that 
Cr0 2 groups are joined by oxygen, thus: 

CrO a <° K 

CrO„<X 

Cr0 »<OK. 



Perchromic Acid. — When hydrogen dioxide is added to a solution of 
chromic acid containing hydrochloric, sulphuric, or nitric acid, an 
indigo-blue solution results, which is supposed to contain perchromic 
acid. The color rapidly disappears, oxygen is evolved, and at the same 
time a chromic salt is formed. It has been found that for every mole- 
cule of potassium dichromate five molecules of hydrogen dioxide are 
required, and eight atoms of oxygen are set free. Based on these facts 
are the following equations: 

K 2 Gr 2 7 + H 2 2 + 2HCl = 2KC1 + 2H 2 + Cr 2 7 . 
Cr 2 7 + 6HC1 + 4H 2 2 = 2CrCl 3 + 7H 2 + 80. 

Cr 2 7 is the anhydride of perchromic acid, HCr0 4 , an analogue of 
HCIO4. 

Perchromic acid is removed from its aqueous solution by ether. The 
ethereal solution of the acid is more permanent than the aqueous solu- 
tion, losing oxygen, however, on evaporation and leaving chromium 
trioxide. 

OK 

Potassium Chlorochromate, Cr0 2 < cl , is obtained by gently heating a 

mixture of 3 parts of potassium dichromate and 4 parts of strong hydro- 
chloric acid and a little water : 

K 2 Cr 2 7 +2IICl = 2Cr0 2 <£K + H 2 0. 

The salt crystallizes on cooling in reddish-yellow crystals. Potassium 
chlorochromate is decomposed by water with formation of potassium 



184 THE SIXTH GROUP. 

dicliromate and hydrochloric acid. It may, however, be recrystallized 
unchanged from aqueous hydrochloric acid. The dry salt is decom- 
posed at 100° as follows . 



4Cr0 2 <^ 1 K = K 2 Cr 2 7 + Crt,0 3 + 2KC14-Cl 2 + 



CI 
Chromyl Chloride or Chromium Oxychloride, CrO^p,* — 

To prepare this compound 10 parts of sodium chloride and 12 
parts of potassium dichloride are fused together. The fused 
mass is broken into lumps and placed in a retort, and 30 parts of 
fuming sulphuric acid are added. The mixture is then heated, 
and the distillate is collected in a cooled receiver. Chromyl 
chloride is a Mood-red liquid, boiling at 118°. Its observed 
gas density corresponds to the formula Cr0 2 Cl 2 . Chromyl 
chloride is decomposed by water, with formation of chromic 
acid and hydrochloric acid. 

Exp. 123.— Dissolve in a flask 20 grams of potassium dichromate in 
150 cc. of water to which 20 cc. of sulphuric acid have been added, a To 
a small portion of the solution add ammonia in excess, and note any 
changes, b. To the remainder of the solution add 10 cc. of alcohol, and 
warm gently. The irritating vapor given off is aldehyde, a product of 
the oxidation of alcohol, c. After the reaction appears to have ceased 
take a portion of b in a test-tube and add ammonia in excess. Note 
result and also the changes which alcohol has produced. Leave the re- 
mainder of b in a covered glass jar for a week or longer. Crystals of 
chrome alum will form. 

Exp. 124. — a. To 100 cc. of a cold saturated solution of potassium di- 
chromate add slowly and with constant stirring 150 cc. of strong sul- 
phuric acid. The solution will deposit on cooling a portion of the 
chromium trioxide formed in the reaction. Pour off the mother-liquor. 
b. Transfer some of the trioxide to a porcelain dish, and add to it a few 
drops of alcohol, c. To a small portion of the trioxide in a test-tube 
add hydrochloric acid and warm the mixture. Chromium trichloride 
will be formed in the solution and chlorine evolved. 

Exp. 125. — Mix 10 grams of pulverized potassium dichromate with 
10 cc. of concentrated hydrochloric acid and 3 cc. of water. Warm 



MOLYBDENUM. 185 

gently until the salt is dissolved. Crytals of potassium chlorochromate 
will form on cooling the solution. 

Exp. 126. — Add a few drops of solution of potassium dichromate and 
also a few drops of hydrochloric acid to about 5 cc. of water in a test- 
tube and pour in sufficient ether to form a layer a centimeter in depth. 
Next add a little barium dioxide, which will react with the acid to form 
hydrogen dioxide. Close the tube with the thumb and agitate, but not 
violently. Note the result. 

The reaction serves as a test for chromium, and also for hydrogen 
dioxide. 



Molybdenum, Mo. 

Atomic Weight, 96. Density, 8.6. 

Molybdenum occurs sparingly in nature. Molybdenum trioxide, the 
compound used in making other compounds of molybdenum, is pre- 
pared from molybdenite, MoS 2 , and wulfenite, the native molybdate 
of lead. 

Metallic molybdenum is separated from its oxides or chlorides by 
hydrogen at high temperatures. When an oxide of molybdenum is 
intensely heated by the compound blowpipe in a carbon-lined crucible 
the metal obtained contains 4 to 5 per cent of carbon. It is very infu- 
sible, and harder than topaz. Molybdenum is permanent in air, and ox- 
idizes only at high temperatures. It is not attacked by h} r drochloric, 
hydrofluoric, or dilute sulphuric acid. It is oxidized by nitric acid, 
and is readily soluble in aqua regia. 

Molybdenum alloys with lead, iron, copper, and gold, making them 
more brittle and less fusible. 

The varying valence of mobybdenum is shown by the compounds 
MoCl 2 , M0CI3, MoCl 4 , and MoCl 5 , Mo0 3 , and MoS 4 . 

Molybdenum Pentachloride, MoCl 5 , is obtained by gently heating mo- 
lybdenum in pure chlorine. It is a black mass, which melts at 194" 
and boils at 268°. Its observed gas density corresponds to the formula 
M0CI5. It may be sublimed in chlorine or carbon dioxide Heated to 
250' in hydrogen it is reduced to the trichloride, M0CI3. This ui 



186 THE SIXTH GKOUP. 

chloride, when heated in carbon dioxide, yields the dichloride and 
tetrachloride: 

2MoCl 3 = MoCl 2 + MoCl 4 . 

The molybdenum chloride, MoCl 6 , corresponding to M0O3, has not 
been obtained. 

Four oxides of molybdenum are known, viz.: MoO, Mo 2 3 , Mo 2) 
and M0O3. The last is an acid anhydride ; the others are basic oxides, 
whose corresponding chlorides have been mentioned. 

Molybdenum Trioxide, Mo0 3 , is commonly called molybdic acid. It is 
a white powder which becomes yellow when heated. It fuses at a red 
heat in a closed vessel, and when heated in air sublimes at the tempera- 
ture at which it melts. It is soluble in acids, and, if it has not been 
heated, is slightly soluble in water. The aqueous solution reacts acid 
with litmus. 

Molybdenum trioxide "is the anhydride of the acids Mo0 2 (OH) 2 , 
MoO v OH) 4 , and Mo(OH) 6 . The first is perhaps formed when solutions 
of molybdates are decomposed by acids, the second separates as a yellow 
crust when a solution of ammonium molybdate in nitric acid stands a 
long time. The hexhydroxide perhaps exists in solution. 

There are a number of polyacid molybdates which ma} r be regarded 
as derived from salts of H 2 Mo0 4 by the addition of M0O3, as illustrated 
by the following sodium compounds ; 



Sodium molybdate, . 
Sodium dimolybdate, 
Sodium trimolybdate, 
Sodium tetramolybdate, 
Sodium octoniolybdate, 
Sodium decamolybdate, 



Na 2 Mo0 4 + 2H 2 0. 
Na 2 Mo0 4 .Mo0 3 . 
Ka 2 Mo0 4 .2Mo0 3 -t 7H 2 0. 
Na 2 Mo0 4 .3Mo0 3 + 6H 2 0. 
Na 2 Mo0 4 .7Mo0 3 -j-4H 2 0. 
Na 2 Mo0 4 .9Mo0 3 + 12H 2 0. 



The constitution of these salts is probably similar to that of the poly- 
chromates. 

Ammonium Molybdate, (NH 4 ) 2 Mo0 4 , separates on adding alcohol to a 
solution of molybdenum trioxide in very strong ammonia. The am- 
monia solution of the trioxide leaves on evaporation crystals of the salt 
(NH 4 ) 6 Mo 7 0, 4 +4HoO. 



Phospbomolybdic Acid.— Phosphoric and molybdic acids unite in dif 
ferent proportions to form compound acids. The one richest in molyb 






TUNGSTEN". 187 

die acid is obtained by evaporating an aqueous solution of phosphoric 
and rnolybdic acids. Yellow crystals separate, having the composition 
2H 3 P0 4 .24Mo0 3 -f-58H 2 0, which at 140° lose water of crystallization. 

Ammonium Phosphomolybdate forms a yellow precipitate when ortho- 
phosphoric acid or an orthophosphate is added to a large excess of a 
solution of ammonium molybdate in nitric acid, known as rnolybdic 
solution. Some investigators have found such a preparation to have the 
composition 2[(NH 4 ) 2 P0 4 ].22Mo0 3 -f 12H 2 0. Other results have been 
obtained which correspond to 20 and 24 M0O3 to 2 atoms of phosphorus. 
Ammonium phosphomolybdate is slightly soluble in water, and is in- 
soluble in rnolybdic solution. It is readily soluble in ammonia Molyb- 
dic solution is an indispensable reagent for the separation of phosphoric 
acid. 

Exp. 127. — Add one drop of solution of common sodium phosphate 
to a few cubic centimeters of rnolybdic solution. An excess of phos- 
phate will prevent the formation of a precipitate. 



Tungsten (Wolfram), W. 

Atomic Weight, 184. Density, 19.1. 

Tungsten is not a very abundant element. Wolframite, an iron and 
manganese tungstate, is the common ore of tungsten. Scheelite, 
CaW0 4 , and stolzite, PbW0 4 , are important tungsten minerals. 

Metallic tungsten is prepared by reducing the oxide with carbon or 
hydrogen. The metallic powder thus obtained has a bright-gray hi si re. 
and is sufficiently hard to scratch glass. Tungsten is more infusible' 
than molybdenum, and has only been fused in an atmosphere of 
nitrogen by means of a powerful galvanic current. The metal in form 
of powder is not changed by moist or dry air, and at a red heal bums 
in air to the trioxide. Hot nitric, hydrochloric, and sulphuric acid 
slowly, and aqua regia quickly, convert the metal into trioxide. A 
boiling solution of potassium hydroxide attacks the metal, potassium 
tungstate being formed and hydrogen evolved. Tungsten is said to 
improve the quality of steel, but little use has as yet been made of 



188 THE SIXTH GROUP. 

The varying valence of tungsten is exhibited by the chlorides, WC1 2 , 
WC1 4 , WC1 5 , and WC1 6 . 

Tungsten Hexchloride, WC1 6 , is formed by the direct union of its ele- 
ments. It melts at 275 and boils at 346°. Its gas densit}^ at tempera- 
tures near its boiling point has been found to correspond to WC1 8 . At 
higher temperatures it dissociates into free chlorine and the following 
compound : 

Tungsten Pentachloride, WC1 5 , is obtainel by repeated distillation of 
the hexchloride in an atmosphere of hydrogen. It melts at 248° and 
boils at 275°. 6, and its gas density accords with WC1 5 . 

Tungsten Tetrachloride, WC1 4 , forms a non- volatile residue w T hen the 
higher chlorides are repeatedly distilled in hydrogen. It is decomposed 
by heat into the volatile pentachloride and tungsten dichloride, which 
remains as an amorphous gray powder. 

There are two oxides of tungsten, W0 2 and W0 3 . The first is a 
feeble base and the last is an acid anhydride. 

Tungsten Dioxide, W0 2 , is obtained as a browm powder when the tri- 
oxide is reduced by hydrogen at a dull-red heat. By acting upon the 
trioxide with zinc and dilute hydrochloric acid, copper-red crystals of 
the dioxide are formed. It dissolves in warm hydrochloric and sul- 
phuric acid to a purple-colored solution. 

Tungsten Trioxide, W0 3 , occurs native. It is obtained by igniting 
tungstic acid, and is commonly prepared from wolframite. The pul- 
verized mineral is digested with hydrochloric acid, and the solution 
containing iron and manganese chlorides poured off. The yellow 
product is dissolved in ammonia to separate any undecomposed mineral, 
and the solution is evaporated to dryness. This residue on ignition is 
converted into the trioxide, a bright yellow powder. At 250° in presence 
of hydrogen it is converted into a blue compound, having the compo- 
sition 2W0 3 -f WOo. Tungsten trioxide is insoluble in water and acid, 
but dissolves in ammonia, and alkali hydroxides and carbonates with 
formation of tungstates. 

Tungsten trioxide is the anhydride of the acids W0 2 (OH) 2 , WO(OH) 4 , 
and W(OH) 6 . 

Dimetatungstic Acid, W0 2 (0H) 2 , separates as a yellow powder when 
an alkali tungstate is treated with an excess of hot acid. It is in- 



TTOGSTEN. 189 

soluble in water, but is dissolved by hydrofluoric acid and an excess of 
concentrated hydrochloric acid. The sodium salt, W0 2 (ONa) 2 , is pre- 
pared by fusing a mixture of tungsten trioxide and sodium carbonate. 
It crystallizes with two molecules of water. 

Monometatungstic Acid, W0(0H) 4 , forms a white precipitate when a 
mineral acid is added to a cold dilute solution of an alkali tungstate. 
It is somewhat soluble in water, and reddens litmus. It changes on 
standing, more rapidly when heated, to the yellow dimetatungstic acid. 

Orthotungstic Acid, W(OH) 6 , is perhaps formed when a solution of 
sodium tungstate is treated with sufficient hydrochloric acid to form 
chloride with all the sodium. After dialyzing the solution for removal 
of the sodium chloride the solution yields on evaporation a gummy mass 
which may be heated to 200° without becoming insoluble in water. 

Tetratungstic Acid, H 2 W 4 O l3 . — The constitution of this acid is sup- 
posed to be analogous to that of the polychromic and molybdic acids. It 
is termed metatungstic acid by some writers. It is obtained by adding 
dilute sulphuric acid to a warm concentrated solution of barium tetra- 
tungstate. The solution deposits on evaporation in vacuum small 
crystals of H 2 W 4 0i 3 -f 7H 2 0, which at 100° lose their seven molecules 
of water. The solution of tetratungstic acid is strongly acid. It may 
be boiled without change and evaporated on a water-bath to a syrup, 
but on further heating suddenly changes to the common yellow tungstic 
acid. 

Sodium Tetratungstate, Na 2 W 4 13 -f- 10H 2 0, is prepared by boiling the 
salt Na 2 W0 4 + 2H 2 with tungsten trioxide. It is very soluble in 
water. 

In addition to the sodium salts already described the following have 
been obtained : 

Na 2 W 2 7 + 2H 2 0. 
Na 4 W 3 11 + 7H 2 0. 
Na 4 W 5 17 + llII 2 0. 
Na 6 W,0 24 + l6H 2 0. 
Na 6 W 7 21 + 21H 2 0. 
Na 10 W 12 O 41 +21 or 25 or 28 H 2 0. 

These sodium polytungstates may be viewed as having a constitution 
analogous to the polysulphates and polychromates, but with this differ- 
ence, namely, that some of the salts contain more than two atoms of 




190 THE SIXTH GKOUP. 

basic metal. These latter are formed from ]STa 2 W0 4 and W0 3 in dif- 
ferent proportions. The salt ]Sra 4 W 3 0ii may be graphically represented 
thus: 

NaO >WO >0 

wo 

HaO wo ><> 

NaO >wu 

The salt Nai Wi 2 O 4 i is manufactured by roasting wolframite with 
soda ash, and exhausting the fused mass with hot water. The boiling 
solution is nearly neutralized by hydrochloric acid and left to crystal- 
lize at common temperatures. This salt is used to render fabrics unin- 
flammable. 

Barium Tetratungstate, BaW 4 0i 3 + 9H 2 0, separates in crystals when 
hot saturated solutions of sodium tetratungstate and barium chloride 
are mixed. 

Tungsten Sodium Bronze. — "When a sodium polytungstate is fused 
with tin and the product successively treated with water, hydrochloric 
acid, and sodium hydroxide, fine golden-yellow cubes of tungsten 
sodium bronze remain. The compound is used for bronze powders. 

Tungstic acid forms very complex compounds with phosphoric acid, 
analogous to the phosphomolybdates. 



Uranium^ U. 

Atomic Weight, 239. Density, 18.7. 

Uranium is mostly obtained from the mineral uranite or pitch blende, 
which consists of U 3 8 and impurities. It also occurs in a few other 
minerals. The metal is best prepared by heating uranium tetrachloride 
with sodium and sodium chloride in an iron crucible closed with a 
screw cap. At a red heat it separates as a grayish-black powder, which 
fuses at a white heat to globules. The fused uranium is silver white, 
somewhat malleable, and not quite as hard as steel. It gradually cor- 
rodes in air, becoming first steel-blue, and finally black. Uranium 
powder ignites in air or oxygen at 150° to 170°, and burns brilliantly. 



URANIUM. 191 

The atomic weight of uranium was for a long time regarded as 120. 
The close analogy of its compounds with those of molybdenum and 
tungsten has, however, recently led to the adoption of double this num- 
ber as the true atomic weight. Further, the determination of the 
specific heat of the metal and the gas densities of some of its compounds 
confirms this view 

The different classes of compounds which uranium is capable of 
forming are illustrated by the following: UC1 8 , UC1 4 , UC1 6 . U0 2 , U 2 5 , 
U0 3 , and the little understood U0 4 , U0 4 + 2H 2 0, K 4 UO 8 + 10H a O. 

Uranium Trichloride, UC1 3 , results when the tetrachloride is heated to 
near the temperature of volatilization in hydrogen : 

UC1 4 + H = UCI3 + HCI. 

Uranium Tetrachloride, UC1 4 , is formed by the direct union of its 
elements, and is prepared by passing dry chlorine over a hot mixture 
of uranium oxide and charcoal. Its observed gas density is 192.9, 
theory requiring 185.2 for UC1 4 . 

Uranium Pentachloride, UC1 5 , is formed by the direct addition of 
chlorine to the tetrachloride. The pentachloride begins to dissociate at 
120° into the tetrachloride and chlorine, and at 235° the dissociation is 
complete. 

The oxides of uranium are UO(?), U 2 3 (?), U0 2 , U0 3 , U0 4 , TJ 2 5 , 
and U 3 8 . 

Uranium Dioxide or Uranous Oxide, U0 2 , is obtained by heating a 
higher oxide in hydrogen. It was formerly regarded as metallic 
uranium. It is a basic oxide, forming green salts. The sulphate has 
the composition U(S0 4 ) 2 -|-8H 2 0. 

Uranium Trioxide or Uranic Oxide, U0 3 , is formed when uranyl nitrate 
is heated to 250°. Towards acids it is basic, and forms salts containing 
the bivalent radical uranyl, U0 2 . Fused with alkali carbonates it forms 
uranates. The uranyl salts are yellow. Glass colored yellow by them 
is remarkable for its fluorescence. 

Uranic Acid or Uranyl Hydroxide, U0 2 (0H) a . — This hydroxide does not 
separate on addition of an alkali hydroxide to a solution of* a uranyl sail 
owing to formation of a uranate. When a solution of uranyl nitrate in 
absolute alcohol is evaporated a violent reaction occurs, and a mass re- 



192 THE SIXTH GROUP. 

mains, which, after treatment with water, leaves a compound of the 
composition H 2 U0 4 . 

Uranyl Chloride, U0 2 =C1 2 , is formed by the direct union of uranous 
oxide and chlorine at a red heat. It is soluble in water without decom- 
position. 

Uranyl Nitrate, ^Q 2 ~j!>U0 2 -f 6H 2 0.— This salt is formed when any 

oxide of uranium is dissolved in nitric acid. It is soluble in half its 
weight of water. 

Uranates. — Potassium diuranate, K 2 U 2 7 , and other diuranates are 
analogous to the dichromates. More complex polyuranates are known. 

Uranium Pentoxide, U 2 5 , is formed when any one of the oxides is in- 
tensely ignited in air. It is a black powder used as a black pigment on 
porcelain. 

Green Oxide of Uranium, U 3 8 , is a dark-green powder formed when 
either U0 2 or U0 3 is ignited in air, or better in oxygen. 

Peruranates. — Hydrogen dioxide produces in a solution of uranyl 
nitrate a yellowish-white precipitate, having the composition U0 4 -|- 
2H 2 0, which dissolves in hydrochloric acid with evolution of chlorine. 
On adding uranyl nitrate to a large excess of sulphuric acid containing 
hydrogen dioxide there separates after some time a colorless crystalline 
powder which is U0 4 . Alcohol throws down from a solution containing 
uranyl nitrate, potassium hydroxide, and hydrogen dioxide, a reddish 
yellow precipitate of K 4 U0 8 + 10H 2 O, which loses oxygen and absorbs 
carbon dioxide on exposure to air. 



Summary of the Sixth Group. 

The sixth group comprises two well-defined sub-groups, one 
containing the non-metals oxygen, sulphur, selenium, and tel- 
lurium, and the other the metals chromium, molybdenum, 
tungsten, and uranium. Oxygen and sulphur are non-metallic 
in all respects, while selenium and tellurium possess some 
metallic properties. All the non-metals of the group form 
acids with hydrogen and oxygen, and in no case bases ; they 



, 



Sulphur, . . 


fe 2 Cl 2 


SC1 2 





sen 


— 


Selenium, 


Se 2 Cl 2 








SeCl 4 


— 


Tellurium, 




TeCl 2 





TeCl 4 


— 


Chromium, . 




CrCl 2 


CrCl 3 





— 


Molybdenum, 




MoCl a 


MoCl 3 


M0CI4 


Mocu 


Tungsten, 




WC1 2 





won 


wen 


Uranium, . 







UC1 3 


ucn 


ucn 



SUMMARY. 193 

also form hydrides, but the metals of the group do not com- 
bine with hydrogen. The oxygen halide OCl 2 is analogous in 
formula, but not in properties, to the dichlorides of other 
elements of the group, and there are no analogues of 2 C1 
and 5 I 2 . The following are halides of this group, oxygen 
excepted: 

SI 6 

CrF 6 

WCL 



The halides of sulphur, selenium, and tellurium are not of 
the nature of halide salts, as is evident from their decomposi- 
tion by water with formation of oxy-acids, thus: 

2S 2 C1 2 + 3H 2 = H 3 SO, + 4HC1 + 3S. 
TeCl 4 + 3H 2 = H 2 Te0 3 -f 4H01. 

Chromium, molybdenum, tungsten, and uranium are basic 
in the lower and acidic in the higher halides. Chromium 
trichloride, for example, is changed by potassium hydroxide 
to Cr(OH) 3 , a base which forms salts with acids. The hex- 
fluoride, on the contrary, reacts with water to form chromic 
acid, thus: 

CrF 6 + 4H a O = H 2 Cr0 4 + 6HF. 
Oxides. 

Sulphur, S0O3 

Selenium, 

Tellurium, 

Chromium, Cr a O s 

Molybdenum, . . , . Mo a 3 

Tungsten, 

Uranium, , 



so 2 


so 3 


SeO a 





TeO a 


TeO a 





CrOs 


MoO a 


MoOa 


WOo 


WO.: 


UOa 


UO a 



194 THE SIXTH GROUP. 



Hydroxides. 



Sulphur, SO(OH) 2 S0 2 (OH) 2 

Selenium, SeO(OH) 2 Se0 2 (OH) 2 

Tellurium, TeO(OH) 2 Te0 2 (OH) 2 

Chromium, Cr(OH) 9 Cr(OH) 3 Cr0 2 (OH) 2 

Molybdenum, .... Mo0 2 (OH) 2 

Tungsten, W0 2 (OH) 2 

Uranium, U0 2 (OH) 2 

The hydroxides of sulphur, selenium, and tellurium are 
acids. The hydroxides Cr(OH) 2 and Cr (OH) 3 are bases, while 
Cr0 2 (0H) 2 is an acid. The salts of molybdic acid, Mo0 2 (OH) 2 , 
and tungstic acid, W0 2 (OH) 2 , and the complex molybdates 
and tungstates, are more stable than the compounds in which 
molybdenum and tungsten are basic. The radical uranyl, 
U0 2 , is acidic or basic in character according as it is 
united to basic or acidic radicals, as for example in potassium 

170 / 0-K 

uranate, Tm 2 > , and uranyl sulphate, SO < n > UO . 

UU 2 \0-K U 

Oxygen, with the lowest atomic weight of any member of 
the sixth group, stands apart in most properties from the 
other elements of this decidedly acidic group. Since the 
acidic character of the hydroxides of the group is closely 
related to the number of atoms of oxygen contained in them, 
it appears that oxygen is acidic in character, and is, as its 
name indicates, an acid former. For example, the lower 
hydroxides of chromium are not acids, but Cr0 2 (OH) 2 is, like 
S0 2 (OH) 2 , a strong acid. Moreover, the basic oxides and hy- 
droxides, as a rule, contain fewer atoms of oxygen than oxygen 
acids. 

The properties which an element exhibits in a compound 
depend upon the relations existing between all the elements of 
the compound. Thus oxygen ma}^ be acidic in acids and not 
in bases. In acids containing the same negative element 
the difference between them depends upon their content of 
oxygen. For example, sulphurous and sulphuric acids differ 
by one atom of oxygen. 



THE SECOND GROUP. 

The elements of this group are bivalent basic metals. 
Beryllium, with the lowest atomic weight, stands apart from 
the other members, and in some properties resembles alumi- 
num of the third group. Mercury, with the highest atomic 
weight in the group, also stands apart, and in mercurous 
compounds resembles in its univalent character the elements 
of the first group. 

Calcium, strontium, and barium are closely related, and 
constitute the well-defined sub-group known as the alkali- 
earth metals. Zinc, cadmium, and mercury are heavy metals, 
which present little similarity to the other members of the 
group. Magnesium is intermediate in properties between the 
alkali-earth metals and zinc. 



Beryllium, Be. 

Atomic Weight, 9. Density, 2.1. 

Beryllium is found in a number of minerals, of which the most 
familiar is beryl, a silicate of aluminum and beryllium. Emerald and 
aquamarine are varieties of beryl, which are valued as gems. 

Metallic beryllium is obtained by heating the chloride with sodium. 
It is a malleable white metal, more fusible than silver. It does not 
decompose water at a red heat. 

Beryllium Chloride, BeCL>, is prepared by heating the oxide mixed 
with carbon in a current of chlorine, and also by heating the metal in 
dry hydrochloric acid gas. The gas density of it has been found to bo 
40; theory requires 89.9 for BeCl a . 



196 THE SECOND GEOUP. 

Beryllium Oxide, BeO, is a -white powder obtained by igniting the 
hydroxide. It is insoluble in water, and after intense ignition dissolves 
slowly in acids. 

Beryllium Hydroxide. Be 0H> 2 . — Ammonia produces in a solution of a 
beryllium salt a gelatinous precipitate which, on drying at 100 : , has 
the composition Be(OH) 2 . It is soluble in acids, caustic alkalies, and 
ammonium carbonate. 

Beryllium Sulphate, BeS0 4 . crystallizes with four and seven molecules 
of water. 



Alkali-Earth Metals. 



Calcium: Atomic Weight, 40. Strontium: Atomic Weight, 
87.5. Barium : Atomic Weight, 137. 

The oxides of these metals are called alkali-earths, because 
in many of their properties they stand intermediate between 
the alkalies and the so-called earthy oxides of aluminum and 
iron. Their hydroxides in basic qualities rank next to the 
alkalies, and are sufficiently soluble in water to exhibit alkaline 
taste and deportment with test papers. 

These metals in their analogous compounds show a grada- 
tion in properties related to their difference in atomic weights. 
The solubility of the hydroxides increases from calcium to 
barium, and the solubility of the chlorides and sulphates 
diminishes in the same order. Calcium hydroxide is changed 
to oxide at a red heat, and barium hydroxide loses water only 
on intense ignition. In no case has the gas density of a com- 
pound of the alkali-earth metals been determined: hence the 
received molecular formulas of their compounds rest on 
analogy. One atom of each of these metals replaces two of 
hydrogen, and they are therefore regarded as dyads. Their 
oxides, hydroxides, and chlorides are illustrated by the for- 
mulas Ca = 0, Sr<QH- Ba<£J. 



CALCIUM. 197 

Calcium, Ca. 

Atomic Weight, 40. Density, 1.58. 

Calcium is a very common element, forming about one six- 
teenth of the crust of the earth. It occurs abundantly in 
limestones and many other rocks, and in the bones and shells 
of animals, and is an essential ingredient of soils. Sea, river, 
and spring waters contain salts of calcium. 

Metallic calcium was first isolated by Davy, in 1808, by sub- 
jecting its oxide mixed with mercury to a powerful galvanic 
current. It is more easily prepared by the electrolysis of its 
chloride, and is also obtained by decomposing calcium iodide 
with sodium at a high temperature in a closed vessel. Calcium 
is a yellow metal, malleable and tenacious, and harder than tin. 
It oxidizes readily in moist air, and decomposes water rapidly 
with evolution of hydrogen. 

CI 
Calcium Chloride, CaCl 2 or Ca< cl , is formed when calcium 

oxide, hydroxide, or carbonate is dissolved in hydrochloric 
acid. 

Exp. 128. — a. Dilute 50 cc. of concentrated hydrochloric acid with 
twice its bulk of water, and add crushed marble in excess. When the 
escape of carbon dioxide has nearly ceased, pour in a little chlorine water 
to oxidize iron compounds which may be present, and heat to boiling, 
then add milk of lime to alkaline reaction, and boil a few minutes. 
Filter from the excess of lime which carries with it any iron, aluminum, 
magnesium, silicon, or phosphorus that may have been present in the 
solution. Make the filtrate acid by a few drops of pure hydrochloric 
acid, and evaporate in a porcelain dish to a syrupy consistence. 

b. Pour part of the liquid into a test-tube, and leave to cool and crys- 
tallize, c. Evaporate the remainder to dryness, and heat the residue 
moderately for some time to drive off all the water. 

d Leave some of the dry salt exposed, and note that it soon deli- 
quesces unless the air is very dry. 

Anhydrous calcium chloride made as described in the above 

experiment is an opaque porous mass which is valuable tor 



198 THE SECOXD GKOUP. 

drying gases. It melts at a red heat, and solidifies on cooling 
to a crystalline mass. The fused salt is also used as a drying 
agent, being serviceable for removing water from various 
organic liquids. 

Calcium chloride separates on cooling a concentrated solu- 
tion in crystals containing six molecules of water. A mixture 
of crushed crystals and snow lowers the temperature suffi- 
ciently to freeze mercury. Both the hydrous and anhydrous 
salt deliquesce in moist air, and are very soluble in water. 

Slaked lime dissolves in a boiling solution of calcium chloride, and 
the solution filtered hot deposits on cooling crystals of calcium hydroxy- 
p OH 

chloride, « $0 -f- ?H 2 0. This compound is decomposed by water 
ta< cl 

into CaClo and Ca(OH) 2 . 

Calcium Bromide, CaBr- 2 , and Calcium Iodide, Cal 2 , are soluble deliques- 
cent salts. 

F 

Calcium Fluoride, CaF 2 or Ca<p occurs as fluor spar 

(fluorite), and is the most abundant compound of fluorine. 
It is used as a flux, and in the manufacture of hydrofluoric 
acid. Unlike the other halides of calcium, it is insoluble in 
water. 

Calcium Oxide, CaO or Ca=0, commonly known as lime 
or quick-lime, is made by heating to bright redness calcium 
carbonate (limestone, marble, or oyster-shells) in lime-kilns. 
In the large kilns, arranged for continuous burning, lime- 
stone or shells and fuel are put into the top from time to time, 
and lime removed from the base. The carbonate is decom- 
posed by heat into calcium oxide and carbon dioxide, thus : 

CO<^>Ca = Ca=0 + CJ^. 

The decomposition of the carbonate is facilitated by the 
presence of steam or gases which dilute the carbon dioxide, 



CALCIUM. 199 

since lime heated to redness in pure carbon dioxide absorbs the 
gas with reproduction of the carbonate. Pure lime is white, 
amorphous, and infusible. It slowly absorbs moisture and 
carbon dioxide from the air, and after a time falls to a powder 
of air-slaked lime. 

OH 

Calcium Hydroxide, Ca(0H) 2 or Ca< Tp ^ s commonly 

known as slaked lime. It is formed more or less readily 
when quick-lime is drenched with water : 

Ca=0 + H-0-H = Ca<^ 

Impure lime, especially that made from limestone contain- 
ing clay, slakes very slowly. 

Exp. 129. — a. Place in an iron or wooden vessel a kilo or more of 
"good lime, and add half its weight of water. After some minutes the 
water is taken up by the lime, which becomes hot, and falls to a dry 
white powder, b. Heat a portion of the dry hydroxide to redness in a 
small glass tube, and observe that water is given off. c. Slake a lump 
of lime in a considerable quantity of hot water. The product is milk 
of lime. 

The heat evolved in the slaking of lime is a result of the 
chemical union of lime and water, and also the passage of the 
latter into the solid state. When large quantities of lime are 
rapidly slaked sufficient heat is produced to char wood. 

Calcium hydroxide is sparingly soluble in cold water, and, 
contrary to the general rule, is less soluble in hot water. One 
part of the hydroxide requires for its solution about 700 parts 
of ice-cold water, and twice as much boiling water. 

Exp. 130. — a. Add some milk of lime to water in a bottle, and filter 
the solution obtained; or cork the bottle and allow (he mixture to stand 
until the lime has subsided, when the clear lime water may be siphoned 
off. b. Heat saturated lime water to boiling, c. Taste lime water, and 
note its reaction with test-papers. 



200 THE SECOND GROUP. 

Lime water readily absorbs carbon dioxide from the air, 
with formation of a film or precipitate of calcium carbonate. 
It is used in medicine as a mild alkali. Lime is the cheapest 
base, and hence finds extensive application in the chemical 
industry. It is used in preparing caustic potash and soda 
from their carbonates, for liberating ammonia from its salts, 
for removing carbon dioxide and sulphur from illuminating 
gas, and for neutralizing acids in various processes. But the 
largest consumption of lime is for making the mortar used in 
ordinary brick and stone work, and for plastering walls. Mor- 
tar is a plastic mixture of slaked lime, sand, and water. The 
pasty slaked lime adheres strongly to the surfaces of the grains 
of sand and building materials ; the admixed water in the 
mortar dries out ; then the mortar becomes tougher owing to 
the gradual absorption of carbon dioxide with formation of 
calcium carbonate and liberation of water, according to the 
equation — 

Ca(OH) 2 + C0 2 = CaC0 3 + H 2 0. 
Hence the prolonged dampness of newly- plastered walls. 

Calcium Dioxide, Ca0 2 — When hydrogen dioxide is added in excess 
to lime water, Ca0 2 + 8H 2 is precipitated in small crystals, which 
effloresce in the air, and become anhydrous at 130°. On ignition Ca0 3 
loses half its oxygen. 

Calcium Hypochlorite, Ca(0Cl) 2 , is formed when hypo- 
chlorous acid is neutralized with lime, and also together with 
calcium chloride by acting on milk of lime with chlorine, thus : 

2Ca(OH) 2 + 2C1 2 = Ca(OCl) 2 + CaCl 2 -f 2H 2 0. 

Bleaching Powder or Chloride of Lime. — This well-known 
substance is extensively manufactured by exposing slaked lime 
to chlorine. The gas is absorbed, and the product is a dry 
white powder, which has a chlorine-like odor, due to the 



CALCIUM. 201 

escape of hypo-chlorous acid, which is set free by the action of 

the carbon dioxide of the air. 

Various views have been held regarding the constitution of 

bleaching powder. Formerly it was regarded as a mixture of 

calcium hypochlorite, calcium chloride, and unchanged slaked 

lime. Odling has represented the portion soluble in water by 

CI 
the formula 0a< O p 1 , while others have viewed bleaching 

powder as a mixture of basic calcium hypochlorite and calcium 
chloride, formed thus : 

3Ca(OH) 2 + 201 2 = 2Ca<^ + CaCl 2 + 2H 2 0. 

Schlaeppi assigns to the bleaching substance the formula 
2CaOCl 2 .H 2 0. 

Bleaching powder, treated with cold water, leaves a residue 
of slaked lime, and the solution contains calcium hypochlorite 
and calcium chloride. Moderately dilute sulphuric and hy- 
drochloric acids liberate chlorine : 

0a01 2 + Ca(001) 2 + 4HC1 = 2C1 2 + 2CaCl 2 + 2H 2 0. 
Ca01 2 + Ca(OCl) 2 + 2H 2 S0 4 = 2C1 2 + 2CaS0 4 + 2H 2 0. 

Bleaching powder gives off the oxygen of the hypochlorite 
when heated to redness. When dilute solutions are boiled 
calcium chlorate is formed. Bleaching powder is largely used 
to bleach cotton fabrics and paper, and as a deodorizer and 
disinfectant. 

Exp. 131.— Shake together a teaspoonful of bleaching powder and 
about 200 cc. of water. Wet thoroughly some bits of calico in the 
solution, then put them into very dilute hydrochloric acid, and repeat 
the treatment with the bleaching" powder solution and acid several 
times. Also try bleaching printed paper, ink and pencil marks. 

Exp. 132. — To a spoonful of bleaching powder in a tall cylinder add 
sulphuric acid diluted with its bulk of water. 



202 THE SECOND GROUP. 

Calcium Sulphate, CaS0 4 or S0 2 <Q>Ca, is abundant m 

nature. It occurs free from water as the mineral anhydrite 
and combined with water as gypsum, CaS0 4 -f- 2H 2 0, of which 
the translucent massive variety is alabaster, and the beauti- 
fully transparent crystallized forms are selenite. Hydrous 
calcium sulphate dissolves in 415 parts of water at 0°, in 368 
parts at 38°, and in 451 parts at 99°. It is commonly found 
together with calcium carbonate in natural waters, and adds 
to their " hardness" and to the formation of boiler-scale. 
When gypsum is heated to 100° or 110° it readily loses most 
of its water, but a temperature of 200° to 250° is required to 
render it anhydrous. 

Calcined gyj)sum, commonly known as plaster of Paris, re- 
covers water of crystallization when mixed with water, and 
hardens to a compact mass, expanding somewhat at the same 
time. It is used for making casts, and mixed with slaked lime 
and sand for hard-finishing walls. Plaster of Paris calcined 
at a low temperature sets more quickly than when strongly 
heated. 

Exp. 133. — Mix plaster of Paris with water to the consistency of thick 
cream. Place a coin, slightly greased, on paper, and pour the plaster 
over it. Rub the coin with the finger to displace air babbles, and leave 
the plaster undisturbed until hard. 



Strontium, Sr. 

Atomic Weight, 87.5. Density, 2.5. 

The more common strontium minerals are the carbonate or 
strontianite, SrC0 3 , and the sulphate or celestite, SrS0 4 . Many 
limestones contain traces of strontium. Metallic strontium is 
obtained by the electrolysis of the fused chloride. It has a 






BAEIUM. 203 

yellow color, and is harder than calcium or lead. It oxidizes 
in the air, and decomposes water rapidly. The strontium salts 
color the flame a brilliant crimson. 

Strontium Chloride, SrCl 2 -j- 6H 2 0, is a deliquescent salt, 
somewhat soluble in alcohol. 

Strontium Oxide, SrO, is obtained by igniting the nitrate. 
With water it forms strontium hydroxide, Sr(0H) 2 , which is 
more soluble than calcium hydroxide, and which crystallizes 
with eight molecules of water. The hydroxide is converted 
into oxide only on intense ignition. 

Strontium Dioxide, Sr0 2 , and its hydrate, Sr0 2 -f- 8H 2 0, have 
been prepared. 

Strontium Sulphate, SrS0 4 , separates as a heavy white pow- 
der on addition of sulphuric acid or soluble sulphate to a 
solution of a strontium salt. It is more soluble in water than 
barium sulphate, and less than calcium sulphate. 



Barium, Ba. 

Atomic Weight, 137. Density, 4. 

Barium is more abundant than strontium, and the chief 
source of it is barite or heavy spar, BaS0 4 . It occurs in many 
other minerals. Metallic barium tarnishes on exposure, and 
decomposes water. It is separated by electrolyzing the fused 
chloride. The soluble barium compounds are poisonous. 
The volatile barium compounds impart a yellowish-green 
color to a non-luminous flame. 

CI 
Barium Chloride, BaCl., or Ba<«-., is the cheapest soluble 

salt of barium, and is therefore extensively used in preparing 



204 THE SECOXD GEOUP. 

other barium compounds. It is made by treating either the 
sulphide or the native carbonate with hydrochloric acid. It 
crystallizes from solution as BaCl 2 -\- 2H 2 in rhombic crystals, 
which are permanent in the air. 

Barium Oxide, BaO, is obtained by igniting the nitrate or 
iodate as a grayish-white substance. It combines energetically 
with water to form the hydroxide. 

Exp. 134. — Heat barium nitrate in a porcelain crucible gently at first 
and finally over the blast lamp. When cool, drop upon it a little water. 
The mass will glow. A strongly alkaline solution of the hydroxide will 
be obtained on dissolving the oxide in hot water. 

OH 

Barium Hydroxide, Ba(0H) 2 or Ba< -rj-, is manufactured 

from heavy spar by a somewhat complicated process. It 
crystallizes from water as Ba(OH) 2 + 8H 2 0. It loses seven 
molecules of water in vacuum over sulphuric acid; at a red 
heat the eighth molecule is driven off, and Ba(OH) 2 remains 
as an oily liquid, which solidifies on cooling. Crystals of 
Ba(OH) 2 -j- 8H 2 dissolve in 3 parts of boiling water and 
in 20 parts at 15°. Barium hydroxide is a strong base much 
used in chemical analysis, and is known as caustic baryta, and 
its solution as baryta water. It absorbs carbon dioxide from 
the air with formation of barium carbonate. 

Barium Dioxide, Ba0 2 , is formed when ox} T gen or dry air free 
from carbon dioxide is passed over glowing barium oxide or 
hydroxide. The conversion is not complete, as the product 
always contains unchanged oxide. On adding a solution of 
barium hydroxide to hydrogen dioxide, the hydrated dioxide, 
Ba0 2 + 8H 2 0, separates in crystals. These, on heating to 
130°, are changed into a white powder of pure anhydrous 
barium dioxide. The dioxide is also formed when a mixture 
of barium oxide or hydroxide and potassium chlorate is heated 
to dull redness. The potassium chloride formed is removed 
by water, the dioxide then remaining as a hydrate. 



BAKIUM. 205 

Exp. 135. — Stir commercial barium dioxide in water and slowly add 
hydrochloric acid until nearly all is dissolved. The solution contains 
hydrogen dioxide and the impurities of the barium dioxide. A strong 
solution of barium hydroxide is added, little by little, to precipitate the 
silica and metallic oxides, until a little hydrated barium dioxide is 
formed. The solution is filtered, and barium hydroxide is added as long 
as the hydrated dioxide separates. The crystals may be washed by de- 
cantation, and kept without drying for use in preparing a pure solution 
of hydrogen dioxide. 

Barium Sulphate, BaS0 4 or S0 2 < Q >Ba, falls as a heavy 

white powder on addition of sulphuric acid or a soluble sul- 
phate to a solution of a barium salt. It constitutes the mineral 
barite or heavy spar, which occurs in large transparent or trans- 
lucent colorless crystals. It is very insoluble, one part of the 
salt requiring 400,000 parts of water for its solution. Ground 
heavy spar is largely used to adulterate white lead, to which 
it is inferior as a pigment. The artificial sulphate has better 
covering properties, and is also used in paints. 

Barium Iodate, Ba(I0 3 )2 -f- 2H 2 0, separates when iodine is added to a 
solution of barium hydroxide. It is soluble in about 600 parts of boiling 
water and in 3000 parts of cold water. 

Barium Monosulphide, BaS, is obtained pure by passing hydrogen sul- 
phide over hot baryta. The crude sulphide is manufactured by roasting 
ground heavy spar mixed with coal-dust. Barium sulphide is decom- 
posed by water, thus : 

2BaS + 2H 2 = Ba(HS) 2 -f-Ba(OH) 2 . 

Barium Hydrosulphide, Ba(HSV 2 , is made by saturating a solution of 
barium hydroxide with hydrogen sulphide. The solution, when evapo- 
rated out of contact with air, yields colorless crystals. When a solution 
of the hydrosulphide is boiled with sulphur, a yellow, caustic, and alka- 
line solution of BaSr, is obtained, which yields on evaporation in vacuum 
free sulphur, and red crystals of the tetrasulphide, BaS,. 



206 THE SECOXD GROUP. 

Magnesium, Mg. 

Atomic Weight, 24.4. Density, 1.75. 

Magnesium is not as abundant as calcium, with which it is 
commonly associated. It is found in plants, milk, blood and 
bones, and in sea and river waters. Many mineral springs 
owe their saline character to magnesium sulphate and chloride. 
Magnesium silicates are common in rocks. Serpentine is a 
hydrous silicate, and talc is an anhydrous silicate of mag- 
nesium. Dolomite, composed of magnesium carbonate and 
calcium carbonate, often occurs in rock masses. Magnesium 
is not found free in nature, although the metal is not altered 
by air and moisture, as are the metals of the alkalies and 
alkali- earths. 

The metal may be separated by electrolysis, or better by 
heating its chloride with sodium. 

Magnesium is almost as white as silver, is malleable, melts 
between 700° and 800°, and distils at a higher temperature. 
Distillation is a means of purifying the metal. The air at 
ordinary temperatures acts but slightly on the metal, since the 
oxide and hydroxide formed adhere firmly, thus protecting it 
from further corrosion. Magnesium, in the form of wire, 
ribbon, and filings, burns with a brilliant, dazzling white 
light, remarkable for its actinic rays. A mixture of chlorine 
and hydrogen may be exploded by this light. Magnesium 
lamps are constructed for burning magnesium ribbon, and are 
used for illumination in photography when sunlight is not 
available. The metal when pure does not decompose water 
at 100°. It dissolves with evolution of hydrogen in dilute 
acids, even carbonic acid water, and in solutions of ammonium 
salts. In the latter case a soluble magnesium ammonium salt 
is formed. 

Exp. 136.— Burn magnesium. Boil water in a test-tube containing a 
piece of magnesium ribbon, and observe whether gas is evolved or not. 



MAGNESIUM. 207 

The magnesium of commerce acts slightly on boiling water. Next add 
a little hydrochloric acid. Magnesium chloride will be formed. What 
is the reaction ? 

CI 
Magnesium Chloride, MgCl 2 or Mg< c ,. — This salt is ob- 
tained by neutralizing hydrochloric acid with magnesia alba, 
adding chlorine water to peroxidize any iron present, and 
then digesting for some time with slight excess of magnesia 
alba. The solution yields on evaporation, best over sulphuric 
acid, deliquescent crystals of the hydrous chloride, MgCl, 
-f- 6H 2 0, which is very soluble in water and alcohol. If the 
hydrous salt is heated to expel the water partial decomposition 
takes place, and the residue consists of magnesium chloride 
mixed with some magnesium oxide. The oxygen of the water 
replaces chlorine, thus : 

MgCl 2 + H 2 = MgO + 2HC1. 

If, however, ammonium chloride is added to magnesium 
chloride, the" solution may be evaporated without formation 
of magnesium oxide. The ammonium chloride is driven off by 
gentle ignition, leaving anhydrous magnesium chloride. In 
another method of preparing the anhydrous chloride the 
hydrous salt is dried by heating in a current of hydrochloric 
acid gas. 

Magnesium chloride unites with other chlorides to form 
well-defined double salts, as for example, KCl.MgCl 3 + 6H 9 0, 
NaCLMgCl, + H 2 0, and 2MgCl,CaCl 2 + 12H 2 0. Ammonium 
magnesium chloride, NII 4 Cl.MgCl, -+- GH 2 0, separates in crys- 
tals when a solution of the two salts is evaporated. Ammonia 
does not precipitate solutions of this double salt, and when 
added to a solution of magnesium chloride only a portion of 
the magnesium separates as hydroxide, and the solution con- 
tains the double salt. It' the magnesium hydroxide formed is 
dissolved by hydrochloric acid and ammonia again added, no 



208 THE SECOND GROUP. 

precipitate will form, owing to the presence of sufficient am- 
monium chloride to form the double salt. 



Magnesium Oxide, Magnesia, MgO or Mg=0, is commonly 
made by heating magnesia alba, which loses water and carbon 
dioxide at a dull-red heat. The magnesia thus prepared is a 
very bulky white powder, which becomes more compact on 
intense ignition. Magnesia is used in medicine. It is a 
strong base, and neutralizes acids. 

OH 

Magnesium Hydroxide, Mg(0H) 2 or Mg<QTT. — Magnesia 

which has not been strongly ignited unites with water with 
evolution of some heat, but not nearly as much as in the 
slaking of lime. The hydroxide is also formed when potas- 
sium or sodium hydroxide is added to a solution of a mag- 
nesium salt. It occurs native as the mineral brucite. 



Magnesium Sulphate, MgS0 4 or S0 Q < >Mg, occurs native 

as epsomite, MgS0 4 + 7H 2 0, and kieserite, MgS0 4 + H 2 0. 
The former, knoAvn as Epsom salt, is named from the Epsom 
springs in England, from the water of which the salt was 
formerly manufactured. Magnesium sulphate has been made 
from serpentine, magnesite, and dolomite, but the Stassfurt salt 
mines, in which kieserite is abundant, yield the bulk of the 
salt in market. Magnesium sulphate is very soluble in water, 
100 parts of water dissolving 27.73 parts of the anhydrous salt 
at 0°, and 73.8 parts at 100°. It crystallizes with seven mole- 
cules of water, six of which are given off below 150° ; but the 
seventh molecule requires a temperature of 200° to expel it. 
If a solution of the salt is evaporated at 70° crystals are ob- 
tained with six molecules of water. Epsom salt has a bitter 
taste, is used as a purgative, and in dyeing with aniline colors, 
and in dressing cotton goods. 



zinc. 209 

ZinCj Zn. 

Atomic Weight, 65.4. Density, 6.9 to 7.2. 

The metal was unknown to the ancients, but they made 
brass by melting copper with zinc ores. Zinc was not smelted 
from its ores until the middle of the eighteenth century, 
although it was known as a distinct metal in the seventeenth 
century. The existence of the free metal in nature is doubt- 
ful. The chief ores of zinc are smithsonite, ZnC0 3 ; cala- 
mine, Zn 2 Si0 4 -\- Hfi ; and zinc blende, ZnS. In New Jersey 
the red oxide of zinc and franklinite (an oxide of iron, zinc, 
and manganese) are worked for zinc oxide. Zinc ores are 
first roasted to expel water, aud in case of blends to change 
the zinc sulphide into oxide. The roasted ore, mixed with 
coal dust, is put into a hot fire-clay retort or muffle. The 
reduction of the zinc oxide by carbon only takes place at a 
temperature above the boiling point of zinc. The vapor of 
zinc is received in a suitable condenser connected with the 
retort. 

Zinc is a white metal with a tinge of blue, and is capable of 
a high polish. It melts at 423° and boils at about 1000°. The 
density of cast zinc is 6.9, and of rolled or sheet zinc 7.2. 
Cast zinc breaks with a highly crystalline fracture : small 
pieces may, however, be hammered into bars. Between 100° 
and 150° it is malleable, and may be rolled into plates, which 
if not too thick maybe rolled cold into sheets. At somewhat 
higher temperatures the metal becomes brittle again, and at 
205° may be pulverized. 

The zinc of commerce, commonly called spelter, usually 
contains small quantities of iron, lead, and arsenic, and traces 
of other metals. It dissolves readily in dilute acids. The 
solution is facilitated by impurities above mentioned, which 
form with the zinc galvanic couples. 

Pure zinc is scarcely attacked by acids unless in contact with 



210 THE SECOND GROUP. 

another metal. The caustic alkalies and ammonia dissolve 
zinc with evolution of hydrogen, the reaction taking place 
more rapidly if a little copper or platinum is precipitated on 
the zinc. 

Zinc does not tarnish in dry air ; but in moist air a thin 
coating of basic carbonate forms, which prevents further cor- 
rosion. 

Zinc is used for a variety of purposes where a cheap sheet 
metal is required that will resist atmospheric action. It is 
also used for the so-called galvanizing of iron. Large quan- 
tities of zinc are used in the manufacture of brass, which is 
an alloy of zinc and copper. 

Zinc dust is a mixture of finely divided zinc and zinc oxide. 
It is formed in the smelting of zinc, and is used to remove 
oxygen from organic compounds, 

Exp. 137. — Melt zinc in a crucible, and pour in a thin stream into 
water. If the metal is poured at temperatures little above its fusing 
point the granulated pieces will be thick. If, however, the metal is 
poured when very hot, thin irregular leaflets will be obtained. 

Exp. 138. — Pour molten zinc into a clean dry mortar, and rub vigor- 
ously with the pestle. As soon as the metal becomes pasty it may be 
pulverized by rubbing. 

Exp. 139. — Place a piece of common zinc in pure dilute sulphuric 
acid, and a similar piece of pure zinc in another portion of acid of the 
same strength. Note which dissolves the more rapidly. After half an 
hour add a drop of platinum chloride or copper sulphate to the acid 
with the pure zinc, and note result. 

Exp. 140. — Warm some of the pulverized zinc of Exp. 138 with solu 
tion of potassium or sodium hydroxide, and add a drop of copper sul- 
phate. Hydrogen will be given off slowly and zinc will dissolve. To 
prove the presence of zinc in the solution, add ammonium sulphide 
a white precipitate is evidence of zinc. Repeat the experiment, using 
instead concentrated ammonia water. 

CI 
Zinc Chloride, ZnCl 3 or Zn< cl< — Finely divided zinc ig- 
nites spontaneously in chlorine, forming the chloride. 



zinc. 911 

Anhydrous zinc chloride is a whitish deliquescent mass. It 
melts at a little above 100°, and distils at a high temperature. 
It has a strong attraction for water, and is used to separate 
the elements of water from organic compounds. The observed 
gas density of it corresponds to the formula ZnCl 2 . 

Zinc chloride is made by dissolving zinc or zinc oxide in hy- 
drochloric acid. The solution becomes viscous on evaporation, 
and yields on cooling crystals of Zn01 2 + H 2 0. A high tem- 
perature is required to expel the last portion of the water. At 
the same time some zinc oxide is formed, thus : 

ZnCl 2 + H 2 = ZnO + 2HC1. 

At a red heat anhydrous zinc chloride distils from the 
residue. 

A concent-ated solution of zinc chloride acts as a caustic on 
the skin, and destroys paper and cloth. Dilute solutions of 
the salt are used for disinfecting purposes. The Burnettizing 
process for preserving wood from decay consists in impreg- 
nating it with a solution of zinc in crude hydrochloric acid. 
Basic zinc chlorides or oxychlorides are formed When zinc 
oxide is boiled with solutions of the chloride. When the 
oxide is mixed with a concentrated solution of zinc chloride 
to a thick paste, the mass soon becomes warm, and hardens. 

A mixture of zinc chloride and hydrochloric acid is much 
used in soldering with common solder, an alloy of tin and 
lead. The acid converts any film of metallic oxide present 
into chloride, which dissolves in the zinc chloride, thus 
cleansing the metal so that the molten solder will wet ami 
adhere to it. Solder will not stick to rusty or oxidized 
metallic surfaces. 

Exp. 141. — Dissolve zinc in hydrochloric acid, and evaporate the solu< 
tion to the consistency of syrup Pour a portion into a test-tube, and 
set aside to crystallize. Put some of the viscous chloride upon cloth 
and paper, and note the result. 



212 THE SECOND GROUP. 

Zinc Oxide, ZnO or Zn=0, is formed when the metal burns 
in air. It is manufactured directly from zinc ores by allowing 
the vapor of zinc obtained in the reduction to burn, and col- 
lecting the zinc oxide formed in suitable chambers. It is a 
white powder, largely employed as pigment known as zinc 
white. It does not blacken on exposure to hydrogen sulphide, 
like white lead. It is insoluble in water and soluble in acids. 

OH 

Zinc Hydroxide, Zn(0H) 2 or Zn< Q Tj, separates as a gelati- 
nous precipitate on addition of sodium or potassium hydroxide 
to a solution of a zinc salt. It dissolves readily in excess of the 
caustic alkali. 

Zinc Sulphide, ZnS, occurs as blende. When an alkali sul- 
phide is added to a solution of a zinc salt a white precipitate 
is formed, consisting of hydrous zinc sulphide or zinc hydro- 
sulphide, soluble in dilute hydrochloric and sulphuric acid, 
but insoluble in acetic acid. Hydrogen sulphide precipitates 
zinc sulphide from solutions containing but little free acid. 

Zinc Sulphate, ZnS0 4 or S0 2 < Q > Zn, is made on a large scale 

by roasting zinc blende, exhausting the mass with water, and 
evaporating the solution. The crystals (see Exp. 15) which 
separate have the composition ZnS0 4 + ~H 2 0, and are known 
as white vitriol and as zinc vitriol. The crystals are isomor- 
phous with Epsom salt. If the solution is concentrated at 50° 
the hydrate ZnS0 4 -f 6H 2 separates in crystals which are 
isomorphous with the corresponding magnesium salt. Zinc 
sulphate is very soluble in water. It is used in medicine, 
in dyeing, and for the preparation of other zinc compounds. 



CADMIUM. 213 

Cadmium, Cd. 

Atomic Weight, 112. Molecule, Cd. Density, 8.6. 

Cadmium was discovered independently by Strom ever and 
Hermann in 1817. It is found in a few minerals, and in 
small quantities in many zinc ores. In the smelting of these 
ores the cadmium distils first, and mostly oxidizes in the re- 
ceivers, forming a brown powder of impure oxide. This is 
mixed with charcoal and distilled, and the metallic product 
is again distilled. To free the cadmium thus obtained from 
the last portions of zinc, and also from copper and arsenic, it 
is dissolved in hydrochloric acid, and to the solution an excess 
of ammonium carbonate is added. The cadmium precipitates 
as basic carbonate, and the other metals remain in solution. 
The precipitate leaves on ignition the oxide, from which the 
pure metal is obtained. 

Cadmium is a brilliant white metal, a little harder than tin, 
and very ductile and malleable. It melts at 315°, and boils 
at 860°. Its gas density at 1040° is 56.9 (Deville and Troost), 
which shows that its molecular weight at high temperatures 
is identical with its atomic weight. The metal is used in 
fusible alloys. An amalgam of cadmium and zinc is em- 
ployed in filling teeth. 

Cadmium Chloride, CdCl 2 , is obtained by dissolving the 
metal or oxide in hydrochloric acid. The solution on evap- 
oration yields crystals of CdCl 2 -f- 2H 2 0, which effloresce in 
air. 

Cadmium Iodide, Cdl 2 , is obtained by treating the metal 
with iodine and water. The anhydrous salt crystallizes from 
water in six-sided plates. It is one of the few metal lie iodides 
soluble ill alcohol. It is used in photography. 

Cadmium Oxide, CdO, is obtained by the combustion of the 



214 THE SECOND GKOUP. 

metal as a brown amorphous powder. It remains as a blue-black 
powder consisting of microscopic octahedra when the nitrate 
is ignited. 

Cadmium Hydroxide, Cd(0H) 2 . — Potassium hydroxide pro- 
duces in solutions of cadmium salts a white precipitate of the 
hydroxide, which is insoluble in an excess of the precipitant 
(difference from zinc hydroxide), but is readily soluble in 
ammonia. 

Cadmium Sulphide, CdS, is found crystallized as the mineral 
greenockite. Hydrogen sulphide throws down from a solution 
of a cadmium salt a brilliant yellow precipitate of cadmium 
sulphide, insoluble in dilute acids. It is employed as a yel- 
low pigment. 

Cadmium Sulphate, CdS0 4 , crystallizes with eight molecules 
of water when its aqueous solution evaporates spontaneously. 



Mercury (Hydrargyrum), Hg. 

Atomic Weight, 200. Molecule, Hg. Density, 13.59. 

This element is also called quicksilver. It is found free in 
nature in small quantities. The chief ore of mercury is cin- 
nabar, HgS. The metal is obtained by roasting the ore, 
whereby the sulphur is burned to dioxide, and the mercury is 
converted into vapor which is condensed in suitable chambers. 
In another process the ore is heated with lime, which combines 
with the sulphur, setting free the mercury which distils off. 

Mercury is a silver-white lustrous liquid, having a density 



MERCURY. 215 

of 13,594 at 0°. It forms when frozen a tin-white malleable 
solid, melting at —39°. 4. Its boiling point is 357° at a pres- 
sure of 760 mm. Its vapor is colorless. At ordinary tem- 
perature the metal is slightly volatile, as may be shown by the 
silvering of gold-leaf which has been kept over mercury for 
some weeks. The density of mercury gas has been found to 
be 100, showing that the mercury molecule contains only one 
atom. 

Pure mercury does not lose its lustre in air at ordinary tem- 
perature, but w T hen heated to near its boiling point it slowly 
absorbs oxygen with formation of red mercuric oxide. Ozone, 
chlorine, and hydrogen sulphide quickly attack it. Hydro- 
chloric acid is without action, and nitric acid dissolves it. 
Mercury may be converted into a fine dust when agitated 
with some substance which will coat the particles. 

Mercury is used in thermometers and barometers, and is in- 
dispensable in experiments with gases. Large quantities are 
employed in the extraction of gold and silver from their ores. 
In the ordinary liquid form it is not poisonous, and has been 
swallowed without injury. When finely divided or in vapor 
it is poisonous. "Blue mass" or "blue pill" is made by 
triturating mercury with a fat or chalk until globules are no 
longer perceptible. 

The mercury of commerce is liable to contain small quanti- 
ties of other metals, which oxidize and form a film, which ad- 
heres to the surface of glass and interferes with its use for 
some purposes. The impurities are best removed by distilling 
the mercury with suitable precautions so that globules of the 
impure metal shall not be carried over with the vapor. For 
most uses in chemical and physical laboratories it can be 
sufficiently purified by treatment with nitric acid, which dis- 
solves zinc, lead, and copper if present. The mercury and 
acid are violently shaken together, or the metal is poured in 
1 thin stream into the acid: the object in cither c;isc is to 
expose a large surface of the metal to the acid. 



216 



THE SECOND GROUP. 



B 



D 



The apparatus Fig. 77 is convenient when mercury requires 
frequent cleansing. The tube B is made with several 
small holes in the end over G, so that mercury may 
fall in a spray into G. The tube G is 60 or 70 cm. long, 
and the tube B is of such length that mercury will fill C 
when in use to a depth of 3 or 4 cm. above the stopper 
which holds B. The tube E is similar to G, and about 
half as long. When the apparatus is first set up sufficient 
mercury is poured into A to fill the small tubes B, B, 
and F. C is then filled with dilute nitric acid and E with 
distilled water. The latter is to stop any acid which may 
possibly pass through B with the mercury. The mercury 
to be cleansed is poured into A, and is received dry in the 
jar G. If it is quite impure the operation must be re- 
peated a number of times. 

Mercury may be freed from dust by straining through 
a filter with a pin-hole in the apex. 

Exp. 142. — Shake in a test-tube some mercury and a 
solution of ferric chloride to obtain the mercury in the 
form of dust. AYash with water by decantation. The 
fine particles may be united by boiling with strong 
hydrochloric acid. 

Amalgams are compounds or alloys of mercury 
with other metals. The amalgams with a large 
proportion of mercury are liquid, and often con- 
tain solid amalgams suspended in the mercury. 
The latter can be separated by pressing it through wet leather, 
leaving the solid amalgam, which in some cases possesses a 
definite composition. Amalgams are commonly obtained by 
putting a clean metal in contact with mercury; sometimes the 
metal is placed in a solution of a mercury salt. Sodium and 
mercury unite with great violence, attended with a brilliant 
flash. Potassium combines without producing light. These 
amalgams are white and when containing more than one part 
of alkali metal to seventy of mercury are hard and brittle. 
By heating to -440°, the crystalline compounds HgK 2 and 
HgXa 3 are obtained. Copper, silver, gold, zinc, cadmium, and 
tin unite readily with mercury. 



MERCUROUS COMPOUNDS. 217 

Exp. 143.— Warm mercury in a test-tube and drop in a small piece 
of clean sodium, and when combination ensues add more. If sufficient 
sodium is added the amalgam when cold will be solid and brittle. 
Place some of the amalgam in water. Hydrogen will be slowly evolved, 
the sodium decomposing the water. 

Exp. 144. — a. Place a clean copper coin in a dilute solution of mer- 
cury nitrate. The mercury will precipitate on the copper, and some of 
the latter will dissolve. 

b. Moisten a piece of sheet zinc with hydrochloric acid, then place it 
in contact with a globule of mercury. 

c. Put a little mercury into a strong solution of silver nitrate contain- 
ing a drop of nitric acid. A crystalline amalgam, known as the silver 
tree or arbor dianae, will form rapidly. 

Mercury forms two classes of compounds, viz. , mercurous, 

in which mercury is univalent, as in Hg-Cl ; and mercuric, 

01 
in which it is bivalent, as in Hg< ™ 

Mercurous Chloride, Calomel, HgCl or Hg-Cl, separates as 
a white powder when a soluble chloride is added to a solu- 
tion of mercurous nitrate. It is obtained by subliming an 
intimate mixture of corrosive sublimate and mercury. Calomel 
is tasteless and insoluble in water and alcohol. It is much 
used in medicine. The density of its vapor has been found 
to be 117.6; theory requires for HgOl 117.7. 

The molecular formula Hg 2 Cl 2 or Cl-Hg-Hg-Cl has been 
assigned to mercurous chloride, mercury being viewed as a 
dyad in mercurous as well as in mercuric compounds. It has 
been supposed that mercurous chloride dissociates completely 
when gasified into mercuric chloride and mercury, which 
combine on cooling to form mercurous chloride again. The 
following facts support this hypothesis. When mercurous 
chloride is sublimed the product contains traces of mercuric 
chloride and minute globules of mercury. It has also been 
observed that calomel vapor amalgamates gold, thus showing 
the presence of free mercury in the vapor. The gas density 



218 THE SECOND GROUP. 

of a mixture of equal molecules of HgCl 2 aud Hg is the same 
as the density of a gas composed of molecules of HgCl, thus: 

2 volumes of mercuric chloride gas weigh . 270.8 
2 volumes of mercury gas weigh .... 200 



4 volumes of the mixed gases weigh . . . 470. 8 
1 volume " " "".... 117.7 

It is obvious that the decomposition of mercurous chloride 
by heat does not decide the question whether its molecule is 
Hg 2 Cl 2 or HgCl, for— 

Hg 2 Cl 2 = HgCl 2 + Hg, and 
2HgCr=HgCl 2 +Hg. 

Eecent experiments by Eileti show that mercurous chloride 
can exist as a gas, and that its molecule is HgCl. He found 
that when a mixture of mercurous and mercuric chlorides 
was heated to 400° no dissociation occurred, and that a gilded 
tube held in the gases was not amalgamated. His experiment 
proves that the presence of mercuric chloride prevents even a 
slight decomposition of mercurous chloride when gasified. 

Mercurous Iodide, Hgl or Hg-I, may be prepared by rub- 
bing together 200 parts of mercury and 127 of iodine. It is 
a greenish powder, very slightly soluble in water, and which 
slowly decomposes into mercuric iodide and mercury. 

Mercurous Fluoride, HgF or Hg-F, differs from the other 
mercurous halides in that it is soluble in water. 

Mercurous Oxide, Hg,0 or HgO — HgO, is prepared by adding 
potassium hydroxide to a solution of a mercurous salt. It is 
a black powder that decomposes in light, and on warming 
into mercury and mercuric oxide. 

Mercurous Sulphate, Hg 2 S0 4 or S0 2 <Q~-rr^, falls as a 
powder when sulphuric acid is added to a solution of a mer- 



MERCURIC COMPOUNDS. 219 

curous salt. It is also formed when sulphuric acid is heated 
with an excess of mercury. 

CI 
Mercuric Chloride, Corrosive Sublimate, HgCl 2 or Hg< cl , 

is formed by the direct union of mercury and chlorine, and 
also when mercury is dissolved in aqua regia. It is manufac- 
tured by heating a mixture of mercuric sulphate and common 
salt, which react thus: 

SO J <°>Hg + 2Na-Cl = Hg.<g+80,<g:g. 

Mercuric chloride dissolves in 13.5 parts of water at 20°, 
and in 1.85 parts at 100°. It is more soluble in alcohol and 
ether than in water at ordinary temperature. It melts at 
about 265°, and boils at 293°, giving rise to a colorless gas 
having a density of 135.4, and corresponding to the molecular 
formula Hg01 2 . It separates from solution in colorless crys- 
tals. When sublimed it forms a translucent crystalline mass. 
Corrosive sublimate is a violent poison. It is used in medicine, 
and for preserving skins, dried plants, and anatomical prepa- 
rations. Keducing agents precipitate mercurous chloride from 
a solution of mercuric chloride, thus: 

2Hg<^j + S0 2 + 2H-0-H = 2Hg-Cl + S0 2 <^+2H-Cl. 

If stannous chloride is used, mercurous chloride separates, 
and is reduced to metallic mercury by an excess of stannous 
chloride : 

2HgCl 2 + SnCl 2 = 2HgCl + SnCl 4 , 
2HgCl + SnCl a = 2Hg + Sn01 4 . 

Mercuric Oxide, HgO or Hg = 0. — As already stated, this 
compound is formed by direct union of its elements. It is 
more easily prepared by heating mercury nitrate, either alone 



220 THE SECOKD GROUP. 

or mixed with mercury; red fumes are given off, and the oxide 
remains as a red crystalline powder. It turns black when 
heated cautiously, the red color reappearing on cooling, and 
at a red heat is completely decomposed. When potassium 
hydroxide is added to a solution of a mercuric salt mercuric 
oxide is obtained as an orange-yellow amorphous precipitate. 

Mercuric Sulphide, HgS or Hg = S. — Mercury and sulphur 
combine when triturated or heated together to form mercuric 
sulphide. Soluble sulphides throw down from solutions of mer- 
curic salts mercuric sulphide as a black precipitate. In case 
mercurous salts are present a mixture of mercuric sulphide and 
mercury results. The black amorphous sulphide changes to 
the red crystalline form when sublimed, and also when heated 
with a solution of an alkali polysulphide. Red mercuric sul- 
phide is a valuable pigment known as yermillion. 

Mercuric Sulphate, HgS0 4 or S0 2 < Q >Hg, is prepared by 

heating mercury with an excess of sulphuric acid as long as 
sulphur dioxide is evolved. It is a white compound, which is 
decomposed by boiling water into the yellow salt Hg 3 S0 6 . 
This is called a basic salt, and as such may be formulated 
thus : 

U< Hg/° 
It may also be viewed as mercuric orthosulphate : 

Exp. 145. — Dissolve mercury in an excess of cold dilute nitric acid of 
density 1.2 to obtain a solution of mercurous nitrate. Dilute the solu- 



MERCURIC COMPOUNDS. 221 

tion largely with water, and in case a yellow precipitate forms, add nitric 
acid until a clear solution results. Read the description of the mercury 
nitrates. To a portion of the solution add hydrochloric acid ; the white 
precipitate is calomel, and will not dissolve in water. To another portion 
add sulphuric acid to obtain mercurous sulphate. Pour some of the 
mercurous nitrate into an excess of solution of potassium hydroxide to 
obtain mercurous oxide. 

Exp. 146. — Dissolve mercury in an excess of hot strong nitric acid. 
To test for the presence of mercurous salt add some of the solution to 
nitric acid diluted with its bulk of water, and then add a few drops of 
hydrochloric acid. In case a precipitate of HgCl forms, add more nitric 
acid to the original solution, and heat again. Pour some of the mer- 
curic nitrate into a solution of potassium hydroxide. The yellow pre- 
cipitate obtained is mercuric oxide. 






THE FIFTH GROUP. 



The members of this group are nitrogen, phosphorus, ar- 
senic, antimony, bismuth, vanadium, niobium, didymium, 



samarium, and tantalum. 



Nitrogen^ N. 

Atomic Weight, 14. Molecule, N 2 . 

Nitrogen exists free in the atmosphere, of which it consti- 
tutes four fifths by volume. Its compounds occur in all 
plants and animals, and are essential components of all soils. 

Nitrogen is a colorless, tasteless, and odorless gas, having a 
density of 14 (hydrogen as unity). It is but slightly soluble, 
100 volumes of water absorbing 1.5 volumes of the gas at 1^°.6. 

Nitrogen liquefies at —146° under a pressure of 35 atmos- 
pheres. Liquid nitrogen has a density of 0.885, and boils at 
— 194°. 4 under a pressure of one atmosphere. It solidifies at 
—214° and 60 mm. pressure. 

Under the influence of electric sparks the pressure of 
nitrogen gas diminishes slightly, according to Thomson and 
Trelfall, who attribute the change to the formation of an allo- 
tropic modification. 

Free nitrogen is remarkable for its chemical inertness 
towards other elements under ordinary conditions. It does 
not combine directly with oxygen when a mixture of the two 



AMMONIA. 223 

gases is heated, and only slowly when subjected to electric 
sparks. As a constituent of the atmosphere, it serves to dilute 
the oxygen, and takes no part in the chemical changes of 
respiration and ordinary combustion. Nitrogen combines 
directly with boron, magnesium, vanadium, and titanium, to 
form nitrides. 

Nitrogen is commonly prepared from air. Burning phos- 
phorus will combine with most of the oxygen, and what re- 
mains may be removed by passing the gas through a tube filled 
with copper turnings heated to redness. Pure nitrogen is 
useful in experiments which must be conducted in an inert 
gas, that is, one which does not react with the substances to 
be tested. Many of the determinations of gas densities at 
high temperatures are made in an apparatus filled with 
nitrogen. 

Exp. 147. — Ignite a dry bit of phosphorus in a shallow porcelain 
dish floating on water, and cover with a bell-jar. The phosphorus 
unites with oxygen to form phosphorus pentoxide, which combines 
with water to- form phosphoric acid, a soluble compound. After the 
cloud lias disappeared introduce a burning splinter or wax taper into 
the gas. 

Ammonia, NH 3 or N(-H. — Nitrogen and hydrogen, under 
-H 

the influence of the silent electrical discharge, combine to form 
ammonia in very small quantity. It is formed in the decay of 
nitrogenous animal and vegetable matters, and also when 
these are subjected to dry distillation, that is, are strongly 
heated out of contact with air. The chief source of ammonia 
and its compounds is the ammoniacal liquor of gas-works. 
Bituminous coal contains about 2 per cent of nitrogen, most 
of which is converted into ammonia in the distillation of the 
coal in the manufacture of gas. Ammonia compounds exist 
in soils, and in minute quantities in the atmosphere. 



224 THE FIFTH GROUP. 

Ammonia is a colorless, pungent gas, which condenses at 
low temperatures to a liquid boiling at —33°. 7, and forms 
when frozen a transparent crystalline body, melting at —75°. 
It dissolves abundantly in water, forming a solution known as 
ammonia, ammonia water, aqua ammonia, and ammonium hy- 
droxide. One gram (1 cc. ) of water at 0° and 760 mm. pressure 
absorbs 0.875 gram or 1144 cc. of ammonia gas; at 10°, 0.679 
gram; at 20°, 0.526 gram; and at 50°, 0.229 gram. Ammonia 
water gives off the gas when open to the air, and all the gas is 
expelled by boiling the solution. 

It may be assumed, for reasons which cannot be given now, 
that ammonia water is a solution of ammonium hydroxide, 
NH 4 0H, formed by the union of NH 3 with H 2 0. Both am- 
monia gas and ammonia water possess strong basic properties, 
neutralizing acids, and forming salts in which the univalent 
radical ammonium, NH 4 , plays the same part as sodium or 
potassium, as illustrated by the following equations : 

NH 3 + H-Cl = NH -01, 
NH 4 -OH + H-Cl = KH 4 -C1 + H 2 0, 
K-OH + H-Cl = K-Cl + H 2 0. 

The first two equations show that the reaction between am- 
monia water and hydrochloric acid can be formulated on the 
hypothesis that ammonia water contains NH 4 0H, and also on 
the assumption that ammonia water is simply a solution of NH 3 . 

The density of ammonia gas is 8.5; its molecular weight is, 
therefore, 17. Experiments have shown that 17 weights of 
ammonia contain 14 weights of nitrogen and 3 weights of 
hydrogen. Hence the molecule of ammonia is represented by 
NH 3 . The gas is slowly but completely decomposed by pass- 
ing electric sparks through it; and it has been found that 2 
volumes of the gas yield when dissociated 1 volume of nitrogen 
and 3 volumes of hydrogen. 



AMMONIA. 



225 



Ammonia does not burn in air unless some combustible gas 
is mixed with it. It may be burned in oxygen. The simpler 
tests for its presence are the odor, the bluing of red litmus 
paper, and the formation of a cloud with acid fumes. 

Ammonia gas for experimental purposes is best obtained by 
heating concentrated ammonia water. In the arts it is set 
free from ammonium chloride or sulphate by means of slaked 
lime: 

2NH -CI + Ca<^| = 2NH 3 + Ca<£|} + 2H 2 0. 



Exp. 148. — Heat concentrated ammonia water in a small flask, and 
pass the gas into an inverted cylinder, as indicated in Fig. 78. The 
hole in the rubber stopper in the 
cylinder should be about twice the 
diameter of the delivery-tube. When 
ammonia escapes abundantly into the 
air, lift the cylinder from the delivery- 
tube and place the stoppered end 
under water. The gas will be quick- 
ly absorbed, and water will fill the 
cylinder. The water will react alka- 
line with litmus. 

Ammonia gas cannot be dried 
by calcium chloride as the two 
combine, but it may be freed 
from moisture by passing it 
through a tube filled with frag- 
ments of potassium hydroxide. 
The gas may be collected over 
mercury. 




Exp. 149.— Neutralize 5 cc. of hydrochloric acid, diluted with its bulk 
of water, with ammonia, and evaporate the solution to dryness in a 
small porcelain dish. The product is ammonium chloride. Will the 
acid alone or the ammonia water leave a residue when evaporated? 



226 



THE FIFTH GROUP. 



Exp. 150. — Dilute 2 cc. of sulphuric acid with thrice its bulk of water, 
neutralize with ammonia, and evaporate the solution to dryness at a 
gentle heat. The product is ammonium sulphate. 
Represent the formation of it by an equation. 

Exp. 151. — a. Triturate in a mortar a little of the 
ammonium chloride, obtained in Exp. 140, with cal- 
cium hydroxide. Observe odor, and test escaping 
gas with a moistened red litmus paper ; also notice 
whether a cloud forms when a rod moistened with 
fuming hydrochloric acid is held over the mixture. 

b. Repeat the experiment, using instead of ammo- 
nium chloride the ammonium sulphate of Exp. 150. 
Represent the reaction by an equation. 

Exp. 152. — Add ammonium chloride to a solution 
of potassium hydroxide, and warm gently. In an- 
other test-tube heat a solution of ammonium sul- 
phate to which potassium hydroxide has been added. 
In the first case potassium chloride is formed, in the 
second potassium sulphate. Formulate the reactions. 

Exp. 153. — Heat strong ammonia water in a small 
flask, Fig. 79, and attempt to burn the ammonia 
gas escaping from tube A. Pass oxygen through B 
into the jacket tube C, and apply a flame to the jet of ammonia. 




Fig. 79. 



Diamide, Hydrazine, N 2 H 4 



N: 

or | 

N: 



:H 2 
:H 9 



-This compound of nitrogen 



and hydrogen was discovered in 1887. It is obtained by decomposing 
an organic compound containing nitrogen. Hydrazine is a colorless gas, 
with a peculiar odor, somewhat like that of ammonia. It is very soluble 
in water, and has a strong alkaline reaction. Hydrazine sulphate, N 2 H io 
H 2 S0 4 , and the hydrochlorate, ]Sr 2 H 4 (HCl) 2 , have been prepared. 

Hydroxylamine, N0H 3 or NH 2 -0H, is formed by the action of nascent 
hydrogen on some of the oxygen compounds of nitrogen. It is known 
only in solution or in combination with acids. It has strong basic prop- 
erties, and, like ammonia and hydrazine, unites directly with acids. 
Hydroxylamine hydrochloride has the composition NH 2 OHHCl. The 
aqueous solution of hydroxylamine decomposes on standing into nitro 
gen, ammonia, and water: 3NH 2 OH = NH3 + 3H2O-I-N3. 






NITROGEN HALIDES. 



227 



Nitrogen Fluoride — When a solution of ammonium fluoride is electro- 
lyzed, oily drops of a highly explosive compound are formed, which is 
supposed to be nitrogen fluoride. 

Nitrogen Chloride, NC1 3 . — This dangerous compound is prepared by 
inverting a thin flask filled with chlorine gas in a warm solution of 
ammonium chloride. The gas is absorbed, and impure nitrogen chloride 
in the form of oily drops separates. These may be collected in a lead 
dish and exploded by touching with a feather moistened with turpentine 
and held on along rod. It may also be obtained by electrolyzing am- 
monium chloride. 

Pure nitrogen chloride is obtained as follows. The product obtained 
by acting on ammonium chloride with chlorine is washed with water 
to remove ammonium chloride, and then exposed for half an hour to 
chlorine. (Gattermann.) 

Nitrogen chloride is one of the most violent explosives. It explodes 
in sunshine, and in contact with oil of turpentine, and when heated to 
95°, and sometimes spontaneously. 

Exp 154. — The apparatus, Fig. 80, consists of a tube A, 3 cm. in 
diameter and 18 cm. long, closed at the bottom with a rubber stopper, 
through which pass two glass 
tubes holding platinum wires, 
which are welded to strips of 
platinum foil in A. The stopper 
is fastened to a block of wood. 
The tube A is filled with a warm 
saturated solution of ammonium 
chloride, and a few drops of oil of 
turpentine are added The wires 
are connected with six Bunsen 
cells. Gas is evolved, and nitro- 
gen chloride rises in minute 
particles, and explodes on com- 
ing in contact with the turpentine 
on the surface of the solution. A 
bell-jar over the apparatus pre- 
vents the turpentine from being 
thrown about. The experiment 
is easily and safely made. 

Nitrogen Iodide, NI 3 , is prepared by adding finely divided iodine 
(precipitated by pouring an alcoholic solution into water) to mosl con- 




Pig. 80. 



228 THE FIFTH GROUP. 

centrated ammonia water at 0°. The mack powder is washed with 
ammonia water, alcohol, and either. Dried in a current of cold air it 
explodes violently It is liable to explode under water The com- 
pound N2I5H results when weaker ammonia is used at common tem- 
perature. It leaves NHI 2 on long contact with water at ordinary tem- 
perature. The black explosive powder commonly prepared by treating 
powdered iodine with ammonia water is a mixture of the above com- 
pounds. 

Exp. 155. — Pulverize a few decigrams of iodine, and pour over it a 
tablespoonf ul of concentrated ammonia water. After a quarter of an 
hour collect the powder on a small filter and wash with cold water. Pin 
the filter with contents to a stick, and leave to dry where the explosion 
will do no harm. The dry black powder explodes when slightly 
jarred, and often spontaneously. The moist powder decomposes vio- 
lently in boiling water 

Ammonium Chloride, NH 4 -C1, has long been known under 
the name of sal-ammoniac. It is formed by the combination 
of equal volumes of hydrogen chloride and ammonia — a strik- 
ing illustration of the production of a solid from the union of 
invisible gases. 

NH 3 + HC1 = NH 4 -C1. 

Exp. 156. — Place a jar of dry ammonia gas over a jar of dry hydrogen 
chloride. The gases will diffuse into each other, and a white powder 
will result. 

Ammonium chloride is manufactured on a large scale by 
saturating hydrochloric acid with ammonia obtained by dis- 
tilling the ammoniacal liquor of gas-works. The solution is 
evaporated and the salt is purified by recrystallization from 
hot water, or it is sublimed. It is also made by heating a 
mixture of common salt and ammonium sulphate, when sal- 
ammoniac sublimes and sodium sulphate remains : 

2Na-Cl + S0 3 <g:||. = 2NH 4 -Cl + SO,<g:g. 



AMMONIUM FLUORIDEo 229 

The sublimed sal-ammoniac of commerce is a tough,- fibrous, 
translucent mass. The salt is obtained as a crystalline meal, 
when a solution is rapidly evaporated. Ammonium chloride 
crystallizes in cubes, octahedrons, and other forms of the regu- 
lar system. At 0°, 28 parts and at 100°, 72.8 parts of the salt 
are soluble in 100 parts of water. A solution of the pure salt 
is neutral to test papers, but on boiling becomes slightly acid 
owing to partial dissociation and escape of ammonia. Am- 
monium chloride cannot exist in the gaseous state, and is 
resolved by heat into ammonia and hydrogen chloride gases, 
which recombine on cooling. When the mixed gases diffuse, 
the ammonia, being less dense than the hydrogen chloride, 
diffuses faster, and may be partially separated. 

Exp. 157. — Heat sal-ammoniac in a test-tube, over the end of which is 
placed a moistened red litmus paper. The test paper will at first be 
colored blue by ammonia, and later acid fumes will redden it. Am- 
monium chloride will collect in the cooler part of the tube. 

The density of the gas obtained by heating ammonium 
chloride has been found to be 13.35, a result which accords 
with the view that the salt does not exist in the gaseous state; 
thus: 

2 volumes of ammonia gas weigh 17. 

2 {i " hydrogen chloride weigh 36.4 



4 " (e the mixed gases weigh 53.4 

1 " " " " " 13.35 

Ammonium chloride is used in the preparation of ammonia, 
in dyeing, and for various other purposes in the arts. It is 
also used in medicine. 

Ammonium Fluoride, NH 4 F, is obtained by subliming a mixture of 
ammonium chloride and sodium fluoride. An aqueous solution may be 
prepared by neutralizing hydrofluoric acid with ammonia. On evapo- 
ration ammonia escapes, and hydrogen ammonium fluoride, ^N B 4 F. BF, 



230 THE FIFTH GROUP. 

crystallizes out. Neither of the salts can be preserved in glass, and 
their solutions etch glass. 

Ammonium Bromide, NH 4 Br, and Ammonium Iodide, NHJ, are similar 
to the chloride. The iodide deliquesces in moist air, and slowly turns 
yellow owing to separation of iodine. 

f)_ fJXT 

Ammonium Sulphate, (NH 4 ) 2 S0 4 or S0 2 <q _^g 4 .— This 

salt, is extensively manufactured by neutralizing sulphuric 
acid with ammonia obtained by heating the ammoniacal 
liquor of gas-works with lime. It forms transparent colorless 
crystals,, isomorphous with those of potassium sulphate. It 
dissolves in less than twice its weight of cold water and in its 
own weight of boiling water. It melts at 140°, and at higher 
temperatures decomposes, giving off ammonia, water, nitro- 
gen, and sulphur dioxide. Ammonium sulphate is employed 
in the preparation of ammonia compounds, and is valuable as 
a fertilizer. 

Ammonium Monosulphide, (NH 4 ) 2 S or NH 4 -S-NH 4 , results 
when two volumes of ammonia gas and one volume of hydro- 
gen sulphide are mixed, and cooled to —18°. It forms color- 
less crystals soluble in water, which decompose at ordinary tem- 
perature into ammonia and hydrogen ammonium sulphide. 
An aqueous solution of the monosulphide is prepared by 
dividing an ammonia solution into two equal parts, satu- 
rating one portion with hydrogen sulphide, and mixing the 
two solutions. 

Hydrogen Ammonium Sulphide, HNH 4 S or H-S-NH 4 . — 

When a mixture of ammonia gas and hydrogen sulphide is 
cooled to 0°, this compound is formed in colorless crystals 
which are volatile at common temperature. A solution of 
hydrogen ammonium sulphide is made by saturating ammo- 
nia water with hydrogen sulphide. The solution gradually 
turus yellow owing to formation of higher sulphides, oxygen 



COMPOUNDS OF AMMONIA WITH METALLIC SALTS. 231 

is slowly absorbed from the air, and ammonium thiosulphate, 
sulphite, and sulphate are formed. On long standing the 
solution becomes colorless again and deposits sulphur, all the 
ammonium sulphide being decomposed. 

The polysulphides (NH 4 ) 2 S 4 , (NH 4 ) 2 S 5 , and (NHJ 2 S 7 have 
been prepared. 

A solution of ammonium sulphide is much used in chemical 
analysis to precipitate certain metals as sulphides, and also to 
dissolve the sulphides of some metals. 

Constitution of Ammonium Salts. — The salts which ammo- 
nia forms with acids are regarded as containing the radical 
NH 4 , since this can be transferred from one compound to 
another, and can be replaced by metals. Further, ammonium 
salts are isomorphous with potassium salts, NH 4 apparently 
having many properties in common with K. As examples 
of analogous compounds may be given NH 4 -C1, K-Ol; 
so .0-NH 4 so O-K 

Exp. 158. — Drop some sodium amalgam into a concentrated solution 
of ammonium chloride. The mercury will swell up and become a 
pasty mass, which is supposed to be an alloy of mercury and ammonium 
formed by the ammonium, NH 4) replacing the sodium of the sodium 
amalgam. In the reaction sodium chloride is formed. 

Ammonium amalgam decomposes rapidly, giving off hydrogen and 
ammonia. 

Compounds of Ammonia with Metallic Salts. — Ammonia 
unites with a large number of salts to form compounds whose 
constitution is not well understood. Only a few of these will 
be noticed. The compound 2Ag01.3]S[H 3 is formed when 
silver chloride is dissolved in ammonia, and also when the dry 
chloride is saturated with dry ammonia. On adding ammonia 
to a solution of copper sulphate a basic salt separates, which 
dissolves in excess of ammonia. The deep-blue solution if 



232 THE FIFTH GROUP. 

concentrated deposits on standing crystals of the compound 
CuS0 4 .4NH 3 + H 2 0. 

Mercurous- Ammonium Chloride, NH 3 HgCl, is a black powder 
formed when precipitated mercurous chloride is exposed to 
dry ammonia. If a solution of ammonia is used, dimercurous- 
ammonium chloride, NH 2 Hg 2 Cl, is formed. Mercuric-ammo- 
nium chloride, NH 2 HgCl, separates as a white precipitate 
when ammonia is added to a solution of mercuric chloride. 



Oxides and Hydroxides of Nitrogen. 

There are five oxides and three oxy-acids of nitrogen: 

Nitrous oxide, N 2 Hyponitrous acid, HNO ? 

Nitric oxide, NO. 

Nitrogen trioxide, N 2 3 . Nitrous acid, NO-OH. 

Nitrogen tetroxide, N 2 4 . 

Nitrogen pentoxide, N 2 5 . Nitric acid, N0 2 -0H. 

Nitric Acid, HN0 3 or N0 2 -0H. — This acid is formed syn- 
thetically in small quantity when electric sparks pass through 
moist air. If dry air in a flask is subjected to a succession of 
sparks from an induction-coil red fumes of the compound 
N0 2 will be seen, and on adding water the color will dis- 
appear; the N0 2 and water reacting, as explained later, to 
form nitric acid. Traces of nitric acid are also found when 
hydrogen is burned in oxygen mixed with a little nitrogen. 
The source of nitric acid is nitrates. These salts are formed 
in the soil by the slow oxidation of nitrogenous organic mat- 
ter in presence of alkali carbonates and calcium carbonate. 
Nitrification takes place best in not too wet a soil and in a 
warm climate. 

Nitric acid is manufactured by heating a mixture of potas- 



NITRIC ACID. 233 

sium or sodium nitrate and sulphuric acid, and condensing 
the acid vapors in a cooled receiver: 

NO a -0-K + SO a <gg = N 2 -OH + S0 2 <^ K . 



The acid thus prepared contains water derived from the 
water contained in the sulphuric acid, and in part from de- 
composition of a portion of the nitric acid by heat. To purify 
the product, it is mixed with its bulk of concentrated sulphuric 
acid and redistilled. The distillate is warmed, and a current 
of dry air passed through it to remove nitrogen tetroxide. 
In this way nitric acid, HN0 3 , containing less than one per 
cent of water has been prepared. It is a colorless fuming 
liquid, which absorbs moisture from air, and is very corrosive. 
Its density at 15° is 1.53. When the pure concentrated acid 
is heated a portion decomposes into nitrogen tetroxide, which 
colors it red or yellow, and water and oxygen. It commences 
to boil at 86°, and the boiling point gradually rises to 120°. 5. 
The residue contains 68 per cent of HN0 3 , distils unchanged, 
and has a density at 15° of 1.414. A more dilute solution 
yields by distillation acid of the same strength. The ordinary 
pure concentrated nitric acid has a density of about 1.4. It 
may be preserved in the dark, but in strong daylight, espe- 
cially in sunshine, it becomes red. A stronger acid known as 
red filming nitric acid consists of HN0 3 , a little water, and 
N0 2 , and has a density of 1.50 to 1.53. It is more energetic 
in its action than the ordinary acid. 

Concentrated nitric acid corrodes and oxidizes many organic 
substances, dissolves sulphur slowly with formation of sul- 
phuric acid, and even oxidizes charcoal. When it acts upon 
certain organic compounds the radical N0 2 is substituted for 
hydrogen. Benzene, C fi H 6 , for example, is converted into 
C fi H & NO,, or C 6 H 4 (N0 9 )' a , according to the strength of acid 
used. Cotton (cellulose), C 12 H 20 O 10 , when placed in a mix- 



234 THE FIFTH GROUP. 

tare of fuming nitric acid and oil of vitriol, does not change 
in appearance, but is converted into gun-cotton or cellulose 
hexnitrate, C 12 H 14 4 (N0 3 ) 6 . The yellow stain which nitric 
acid produces on the skin and on cloth is due to the forma- 
tion of a yellow nitrate or nitro compound. 

Nitric acid is extensively used in the manufacture of metal- 
lic nitrates, aniline dyes, sulphuric acid, gun-cotton, and 
nitro-glycerine. 

The hydroxides of pentavalent nitrogen are, theoretically, 
N(OH) 5 , NO(OH) 3 , and N0 2 OH. Certain basic salts may be 
viewed as derivatives of pentabasic or tribasic nitric acid ; but 
the ordinary nitrates are derivatives of monobasic nitric acid. 
Nitric acid reacts with bases to form nitrates, thus : 

N0 2 -0H + K-0H = N0 2 -0-K + H 2 0. 
2N0 2 -OH + Ca<°| = |^Io>Ca + 2H 2 0. 

Metals dissolve in nitric acid with formation of nitrates, and 
at the same time nitrous fumes are evolved, a part of the acid 
being reduced to the lower oxides of nitrogen by the metal, or 
by the hydrogen which the metal replaces. If the acid is 
dilute and cold it will dissolve zinc, for example, without for- 
mation of red fumes, and the nascent hydrogen will convert 
part of the acid into ammonia, thus : 

HN0 3 + 8H = NH 3 + 3H 2 0. 

The reduction is complete when a small portion of a nitrate 
is acted upon by hydrogen evolved from zinc in a solution of 
potassium or sodium hydroxide. 

Exp. 159. — Place 40 grams of potassium nitrate in a pint retort, and 
add the amount of sulphuric acid required to form HKS0 4 . Calculate 
the weight and number of cubic centimeters required of sulphuric acid, 
density 1.8, and containing 86 per cent of H 2 S0 4 . Support the retort 
on a lamp-stand, and heat cautiously with a lamp as long as acid distils. 



NITRIC ACID. 



235 



The neck of the retort is thrust into a flask which rests in a water-pan. 
and is cooled by a wet towel, or by pouring water over it. Towards the 
close of the distillation red fumes will appear, and color the distillate. 

Exp. 160. — Dilute the nitric acid of the previous experiment with its 
bulk of water, and neutralize with ammonia. Evaporate the solution, 
towards the last without boiling, until a drop taken out on a cold rod 
solidities at once. Label the salt ammonium nitrate, and set aside for a 
future experiment. 

Exp. 161. — a. Drop a few bits of pulverized zinc into concentrated 
nitric acid contained in a test-tube. Note observations. 

b. Mix thoroughly 5 cc. of strong nitric acid with 60 cc. of water. 
Pour this dilute acid into a test-tube containing a few grams of pul- 
verized zinc, and observe that the action is very different from that in 
a. After 15 to 30 minutes pour some of the solution into another test- 
tube, add potassium hydroxide in excess, 
heat, and test escaping vapors for am- 
monia. 

Exp. 162. — Place pulverized zinc, not 
too little, and a solution of potassium hy- 
droxide in a test-tube ; warm gently until 
hydrogen escapes freely. The gas will 
not react for ammonia if the potassium 
hydroxide used is free from nitrates and 
nitrites. Add one drop of nitric acid, 
and in a few minutes test for ammonia. 

Exp. 163. — Place 50 cc. of red fuming 
nitric acid in a beaker, standing on a 
plate, which is partly filled with lime 
water. Cover with a bell- jar, as shown in Fig. 81, and drop into the acid 
a piece of glowing charcoal. The charcoal will burn brilliantly on the 
surface of the acid. 




Fig. 81. 



236 THE FIFTH GROUP. 

Nitrates. 






Potassium Nitrate, KN0 3 or N0 2 -0-K. — This salt is also 
called saltpetre and potash nitre. The formation of nitrates 
in soils has already been noticed. Nitre is obtained from the 
nitrification of urine and manure in soil mixed with wood 
ashes to furnish the potassium. The process is carried out in 
artificial nitre beds, which formerly furnished much of the 
nitre used in Europe. The climate of portions of India is 
favorable to the formation of nitrates, and that country fur- 
nishes a very considerable portion of the potassium nitrate of 
commerce. At the present time potassium nitrate is manu- 
factured from Chili saltpetre (sodium nitrate) and potassium 
chloride. These salts are dissolved in hot water, when they 
mutually decompose as follows : 

K-Cl + N0 2 -0-Na = NO -O-K + Na-Cl. 

The sodium chloride separates from the hot solution and is 
removed, and the potassium nitrate crystallizes on cooling the 
solution. The separation depends upon the difference in 
solubility of the two salts. Potassium nitrate forms large 
rhombic crystals when a solution evaporates slowly, or a hot 
saturated solution cools slowly. A crystalline meal is obtained 
by rapidly boiling down a solution. The salt is very soluble 
in water, 100 parts of water dissolving at 0°, 13.3 parts of 
KN0 3 ; at 20°, 31.2 parts; at 100°, 247 parts; at 114°. 1, 
327.4 parts. The boiling point of the saturated solution 
is 114°. 

Potassium nitrate is a powerful oxidizing agent, yielding its 
oxygen readily to combustible substances. It is used in fire- 
works, in the laboratory, in medicine, in salting meats, but 
chiefly in gunpowder. Gunpowder is an intimate mechanical 
mixture of nitre, sulphur, and charcoal. The explosive effect 
is due to the rapid liberation of gases in the burning of the 



KITEATES. 237 

sulphur and charcoal, at the expense of the oxygen of the 
nitre. 

Sodium Nitrate, Chili Saltpetre, NaN0 3 or NO-0-Na.— 

This nitrate is abundant in southern Peru, Bolivia, and 
northern Chili. It is largely used, as already stated, in the 
manufacture of potassium nitrate, and as a source of nitric 
acid. It is a valuable fertilizer. It closely resembles potas- 
sium nitrate in properties, but cannot be substituted for the 
latter in gunpowder, since it becomes moist in air. 

One hundred parts of water at 0° dissolve 73 parts ; at 40°, 
102 parts ; and at 100°, 180 parts of JSTaN0 3 . The saturated 
solution boils at 120°, and contains 100 parts of water and 
216 parts of the salt. 

Ammonium Nitrate, NH 4 N0 3 or N0 2 -0-NH 4 , is obtained by 
neutralizing nitric acid with ammonia or ammonium carbonate 
(see Exp. 160). The salt is decomposed by heat, as stated 
under nitrous oxide, p. 242. It is soluble in hali its weight of 
water at common temperature, and is much more soluble in 
hot water. 

Cupric or Copper Nitrate, Cu(N0 3 ) 2 or xO-0 >Cu— Tllis 
salt is prepared by dissolving copper or copper oxide in nitric 
acid. The solution yields on evaporation deliquescent crystals 
of Cu(N0 3 ) 2 -f 3H 2 0, which are soluble in alcohol. The salt 
loses water and acid at a gentle heat, and is converted into a 
basic salt, and at a red heat cupric oxide remains. 

Silver Nitrate, AgN0 3 or N0 2 -0-Ag. — If a piece of silver 
is hung in cold nitric acid of density 1.2, the metal slowly dis- 
solves without effervescence, the liquid becoming blue from the 
nitrous acid formed. If the acid is hot, the metal dissolves 
rapidly, and nitric oxide is evolved. The salt separates in 
large transparent crystals on evaporating and cooling the 
solution. 



238 THE FIFTH GEOUP. 

According to Kremers, 100 parts of water at 0° dissolve 
121.9 ;' at 54°, 500 ; and at 110°, 1111.0 parts of silver nitrate. 
It is also soluble in alcohol. Silver nitrate melts at 198°, and 
solidifies on cooling to a white fibrous mass. At a red heat it 
is completely decomposed. Its aqueous solution is neutral, 
and has a bitter metallic taste. It is used in photography 
and chemical laboratories, and in small doses as a medicine. 
It acts as a powerful caustic on the skin, turning it black, and 
destroying it. For use as a caustic it is cast into small sticks, 
and is called lunar caustic. Pure silver nitrate is not changed 
by light, Unless in contact with organic matter, when it is 
decomposed, and a dark substance is formed. Indelible inks 
are made of silver nitrate and gum-arabic. The mark on the 
cloth turns black in sunshine or on warming. Better indel- 
ible inks are made of coal-tar blacks, which are not bleached 
by chlorine or potassium cyanide. 

Auryl Nitrate, Au0.N0 3 or N0 2 -0-Au=0, is prepared by 
digesting auryl hydroxide with nitric acid of density 1.40 on a 
water-bath. The solution is evaporated over caustic alkalies, 
which absorb the acid. By this method Schrottlaender ob- 
tained black crystals having the composition 5AuOX0 3 -f-H 2 0. 

Acid Auric Nitrate, HN0 3 . Au(N0 3 ) 3 + 3H 2 0, is obtained by 
treating auryl hydroxide with 3.6 parts of pure nitric acid of 
density 1.49, and heating gently for some hours on a water- 
bath. The yellow solution, after the particles of reduced gold 
have subsided, is decanted and cooled with ice and salt, or is 
evaporated somewhat at 60°-80°, and finally over lime and 
caustic soda. Large crystals separate, which may be kept un- 
changed in a tight-stoppered jar. The salt is decomposed by 
water into auryl hydroxide, AuO.OH. 

Calcium Nitrate, Ca(N0 3 ) 2 or NQ -~Q>Ca. — This salt is 
prepared by neutralizing nitric acid with lime or calcium car- 



NITRATES. 239 

bonate. It is the most abundant product of nitrification in 
soils, but because of its solubility and deliquescence in moist 
air it does not often appear in nature in the solid state. 

Strontium Nitrate, Sr(N0 3 ) 2 or ^ 2 ~Q>Sr. — This com- 
pound is used in pyrotechny for imparting a red color to 
flames. It crystallizes with four molecules of water, which it 
loses at 100°. 

Barium Nitrate, Ba(N0 3 ) 2 or -m-() 2 -0 > ^ a * — ^his sa ^ * s P re " 
pared by dissolving barium carbonate in nitric acid, and also 
by mixing concentrated hot solutions of barium chloride and 
sodium nitrate. Barium nitrate separates when the solution 
cools, and sodium chloride remains in solution. Barium 
nitrate forms anhydrous crystals, which are permanent in the 
air and are sparingly soluble in water. It is used as a re- 
agent, and for making green lights. 

Magnesium Nitrate, Mg(N0 3 ), or N Q 2 ~5>Mg— This salt 



is prepared by neutralizing nitric acid with magnesia alba. 
On evaporating the solution crystals separate, having the com- 
position Mg(NO s ) a -f- 6H 2 0, and which deliquesce completely 
in moist air. The salt is soluble in alcohol. 

Zinc Nitrate, Zn(N0 3 ) 2 or nO-O^ 11 — ^ ns * s a deliques- 
cent salt, which crystallizes with six molecules of water, and 
which is soluble in alcohol. 



Cadmium Nitrate, Cd(N0 3 ) 2 or jJo_o>Cd, crystallizes with 

four molecules of water. It is deliquescent, and soluble in 
alcohol. 



240 THE FIFTH GROUP. 

Mercurous Nitrate, HgN0 3 or N0 2 -0-Hg. — This salt is ob- 
tained by acting on mercury with cold nitric acid of density 
1.2. After a time crystals having the composition HgN0 3 + 
H 2 separate. The salt is soluble in a small quantity of warm 
water, but is decomposed by much water with formation of 
an insoluble basic salt. The presence of nitric acid prevents 
this decomposition/ and to obtain a dilute solution of the salt 
the concentrated solution is added to dilute nitric acid. A 
solution of mercurous nitrate stains the skin at first purple 
and then black. 

Mercuric Nitrate, Hg(N0 3 ) 2 or jto-0 >H ^' — ^- solution of 
this salt is obtained by dissolving mercury in nitric acid, and 
heating until calomel does not separate on adding a drop of 
the solution to dilute hydrochloric acid. The solution yields 
on ]ong standing over oil of vitriol crystals of 2Hg(N0 3 ) 2 -f- 
H 2 0, and the syrupy mother-liquor has the composition cor- 
responding to Hg(X0 3 )„ -[- 2H 2 0. Both the crystals and solu- 
tion are decomposed by water ; at first a white basic salt 
separates, this becomes reddish, and if the washing with water 
is continued mercuric oxide remains. The white salt has the 
composition 3HgO.N 2 5 + H 2 0. 



Nitrous Acid, HN0 2 or N0-0H. — This acid has not been 
prepared in the pure state, but is supposed to be formed 
when nitrogen trioxide is passed into ice-water. A blue solu- 
tion is obtained, which on warming gives off nitric oxide, 
nitric acid at the same time being formed : 

3NO-OH = N0 2 -OH + 2X0 + H 2 0. 

The salts of nitrous acid are stable and important. They are 
distinguished from nitrates by giving off red fumes when 
treated with acids. 



KITROUS OXIDE. 241 

Potassium Nitrite, KN0 2 or NO-O-K, is a very soluble and 
deliquescent salt. It is best prepared pure by the mutual 
decomposition ol silver nitrite and potassium chloride, thus : 

NO-O-Ag + K-Cl. = NO-O-K + Ag-Cl. 

The silver chloride is removed by nitration, and the solution 
is evaporated in vacuum. Potassium nitrate when heated 
gives oft' oxygen, and is more or less completely changed to 
nitrite. The addition of lead to combine with the oxygen 
facilitates the conversion of the nitrate into nitrite. 

Silver Nitrite, AgN0 2 or NO-0-Ag, is prepared by mixing 
hot concentrated solutions of potassium nitrite and silver 
nitrate. Silver nitrite separates as the solution cools. It is 
sparingly soluble in water. 

Hyponitrous Acid. — By treating a solution of potassium nitrate with 
sodium amalgam, neutralizing the solution with acetic acid, and then add- 
ing silver nitrate, a precipitate of silver byponitrite is obtained. Some 
chemists regard this salt as having the composition Ag 2 N 2 0o, and the 
results obtained by others correspond to the formula Ag 4 N 4 5 . The 
free acid is obtained in solution by decomposing the silver salt w r ith 
dilute hydrochloric acid. It decomposes on boiling with evolution of 
nitrous oxide. 

Nitrous Oxide, N 2 0, is a colorless gas which condenses at 
— 88° under ordinary atmospheric pressure to a mobile liquid 
solidifying at about —100°. Liquid nitrous oxide is kept 
in iron flasks, and is a commercial article. The gas has 
a slight odor and sweet taste. At 0°, 100 volumes of water 
absorb 130 volumes, and at 20°, G7 volumes of the gas. Ni- 
trous oxide has been found to contain 28 weights of nitrogen 
and 16 weights of oxygen, and to have a gas density o( 22. 
These data show that the molecule is N 2 0. When the oxygen 
is removed from two volumes of the gas two volumes of nitro- 
gen are left. Nitrous oxide is the anhydride of hyponitrous 



242 THE FIFTH GROUP. 

acid, and is probably produced when the latter decomposes, 
thus : 

H.N.O, = X 2 + H 2 0. 

The oxide, however, does not unite with water to form an 
acid. 

Nitrous oxide gas supports combustion, and most substances 
which burn in air burn with a brilliancy in the gas, recalling 
combustions in oxygen. A glowing splinter inflames in ni- 
trous oxide, and phosphorus burns as in oxygen. Burning 
sulphur is extinguished in the gas unless heated almost to 
boiling. Equal volumes of hydrogen and nitrous oxide ex- 
plode violently when ignited. 

Nitrous oxide gas is used as an anaesthetic in dentistry and 
short surgical operations, where insensibility for a brief period 
only is required. When a mixture of air or oxygen and nitrous 
oxide is breathed for a short time a nervous excitement, often 
accompanied by laughter, is produced, and without loss of con- 
sciousness. Hence the name laughing-gas has been given 
to nitrous oxide. When the gas is to be inhaled it should be 
perfectly pure, and free from chlorine and nitric oxide. 

Nitrous oxide is prepared by heating ammonium nitrate, 
part of the salt being decomposed thus : 

2 X-0-XH 4 = X-O-X + 2H 2 0. 

At the same time a portion dissociates into nitric acid and 
ammonia, which unite in the cooler part of the apparatus to 
form ammonium nitrate again : 

NO.-0-XH 4 = X0 2 -0H + XH 3 . 

When ammonium nitrate is heated too rapidly it decomposes 
with explosive violence and formation of other products. 
Nitrous oxide is freed from nitric oxide by contact with a 
solution of ferrous sulphate, and from acid fumes by potas- 
sium hydroxide. 



NITROUS OXIDE. 



243 



Exp. 164.— The flask A, Fig. 82, has a capacity of about 200 cc. Place 
in it the ammonium nitrate of Exp. 160, and 
in B a piece of red litmus paper. Heat A 
cautiously, so as to avoid too rapid evolution 
of gas, and collect the gas over water. Ob- 
serve that the test paper in B at first turns 
blue and later becomes red. Why ? Do not - 
decompose quite all the salt in A, and when 
through heating, remove the delivery-tube 
from the water. The bottle B will contain 
water and ammonium nitrate. The presence 
of the latter may be made evident by evapo- 
rating the solution cautiously. 



Exp. 165. — a. Thrust a glowing splinter 
into nitrous oxide. 

b. Introduce into the gas on a chalk spoon 
a bit of sulphur burning feebly. 

c. By means of a lamp flame heat the sul- 
phur so that it will burn in the air very 
rapidly, and then thrust it into nitrous oxide, 
enough it will burn in the gas. 




Fig. 82. 



If the sulphur is hot 



Exp. 166. — To show the volume of nitrogen remaining after the oxygen 
has been removed from nitrous oxide (and from nitric oxide) the appara- 
tus, Fig. 83, devised by Prof. E. H. Keiser* may be used. The experi- 
ment is based upon the fact that glowing copper combines with the 
oxygen of the oxides of nitrogen, and leaves the nitrogen free. The 
following is Prof. Reiser's description of the apparatus: 

A represents a gas burette for measuring the volume of the gas. B is 
a gas pipette, which is filled with water The connecting-tube C is made 
of hard glass, and is 3 mm. internal diameter and from 10 to 12 cm. 
long. It is completely filled with granular metallic copper, which has 
been obtained in the reduction of the granular oxide in a current of 
hydrogen. The copper is held in place by plugs of asbestus at each end 
of the tube. To decompose the oxides of nitrogen this tube is heated to 
a red heat with a Bunsen burner, and to prevent it from bending, a piece 
of wire gauze is wrapped around the outside and secured by wives. 

A measured quantity of nitrous or nitric oxide contained in the burette 
A is passed over, the heated copper by opening the pinch-cock and rais- 



Amcricaii Chemical Journal, viii. 92. 



244 



THE FIFTH GKOtXP. 



ing the reservoir tube of the burette. When the reservoir tube is low- 
ered the gas is drawn back from the pipette. It now consists of nitrogen, 
and its volume may be readily determined. If nitric oxide has been 




Fig. 83. 



used the volume of the nitrogen will be exactly one half the original 
volume of the gas; while in case of nitrous oxide the volume of nitrogen 
will be the same as the volume of nitrous oxide taken. 



Nitric Oxide, NO, is a colorless gas, which is distinguished 
from all other colorless gases by forming red fumes of nitrogen 
tetroxide when mixed with oxygen or air. As these fumes 
are irritating, the odor and taste of pure nitric oxide is un- 
known. It is less easily liquefied than nitrous oxide, and, 



NITRIC OXIDE. 245 

like all difficultly condensable gases, is but slightly soluble in 
water. Its gas density is 15, and it contains equal volumes of 
oxygen and nitrogen. Nitric oxide contains the same amount 
of oxygen as a like bulk of nitrous oxide, but does not sug- 
port ordinary combustion, and the affinity between nitrogen 
and oxygen in NO appears to be greater than in N 2 0. A hy- 
drogen flame, burning sulphur, and a candle flame, are ex- 
tinguished when placed in nitric oxide, and so is phosphorus 
when feebly burning, but when burning brilliantly in air it 
burns with increased brilliancy when pluuged into nitric 
oxide. A mixture of nitric oxide and vapor of carbon clisul- 
phide burns with a dazzling light remarkable for its actinic 
rays. 

Nitric oxide is commonly formed when metals are dissolved 
in nitric acid. It is prepared by acting on copper with nitric 
acid. Probably the change takes place in two stages, thus : 

2N0 2 -OH + Cu = ^Jo -0 >Cu + m > 
. NO -OH + 3H = NO + 2H 2 0. 

Thus prepared, the gas is liable to contain nitrous oxide, from 
which it can be freed by passing the gas into a solution of 
ferrous sulphate which absorbs the nitric oxide. The solution 
when heated gives off pure nitric oxide. 

Exp. 167. — Place some copper clippings in a generator (Fig. 52, p. 54), 
and pour in water until the end of the thistle-tube is covered; then add 
concentrated nitric acid, and from time to time more acid as required. 
Collect several bottles of the gas over water. 

Exp. 168. — a. Lift a bottle of nitric oxide from the water pan, and 
place in the red fumes a blue litmus paper. For explanation of acid 
reaction see Nitrogen Tetroxide. 

b. Introduce a burning splinter, and a bit of burning sulphur on a 
chalk spoon into the gas. 

c. Ignite a bit of phosphorus held just above a jar of nitric oxide, 
and quickly plunge into the gas. It will be extinguished. Brine the 



246 THE FIFTH GROUP. 

phosphorus into the air again, and allow it to burn more brilliantly than 
before; then put it into the gas. 

Exp. 169. — Fill a tall cylinder with nitric oxide, cover the end of the 
cylinder with a ground-glass plate, and place upright on the table. 
Drop into the cylinder a thin glass bulb, containing a few cubic centi- 
meters of carbon disulphide. Shake violently to mix the carbon disul- 
phide vapor with the gas ; remove the glass plate and bring a lamp 
flame near the mouth of the cylinder. If the gas has been collected 
over cold water it should be allowed to stand until it has the tempera- 
ture of the room, as the carbon disulphide will not vaporize sufficiently 
at low temperatures. 

Nitrogen Trioxide or Nitrous Anhydride, N 2 3 . — When 
starch is heated with nitric acid of density 1.3 to 1.35, red 
fumes are evolved which condense in a tube surrounded by a 
freezing mixture to a green liquid containing nitrogen triox- 
ide and tetroxide. If nitric oxide is passed into the warmed 
liquid, and the vapors passed through a hot tube, and again 
cooled, pure nitrogen trioxide will be obtained. At —10° it is 
an indigo-blue liquid, which gives off brownish-red vapors 
containing nitric oxide and nitrogen tetroxide. 

Nitrogen Tetroxide, N 2 4 , or Nitrogen Dioxide, N0 2 , is 

formed when one volume of oxygen is mixed with two volumes 
of nitric oxide : 

NO + = N0 2 . 

It is best prepared by heating lead nitrate, which decomposes 
as follows : 



NO.-O 



NO.-O 



>Pb = Pb=0 + 2M) 2 + 0. 



Pure nitrogen tetroxide is a liquid which solidifies to 
colorless crystals at —9°. At a few degrees above its melting 
point the liquid is colorless, but at 15° is orange-colored. It 
boils at 26°. 7, giving off brownish vapors, which become 
darker, and finally, at a higher temperature, almost black. 



NITROSYL CHLORIDE. 247 

At low temperatures the density of the vapor has been found 
to approach that required by the formula N 2 4 . The density 
diminishes with rising temperature, and at 140° corresponds 
to N 2 . These results show that the molecules of N 2 4 are 
dissociated by heat into molecules of N0 2 . Nitrogen dioxide 
and water react at ordinary temperature to form nitric acid 
and nitric oxide, thus : 

3N0 2 + H 2 = 2HN0 3 + NO. 

This explains the acid reaction of the red fumes formed 
when nitric oxide mixes with air. 

Exp. 170. — Place in a bell- jar over water one measure of nitric oxide, 
and quickly mix with it one measure of oxygen. The gases will com- 
bine, and at first expand owing to the heat evolved ; then the red fumes 
will be rapidly absorbed by the water, and a small bulk of nitric oxide 
will remain, as may be proved by adding more oxygen. If the experi- 
ment is made with pure nitric oxide, and oxygen added slowly, nearly 
all the gas in the bell- jar will disappear. 

Nitrogen Pentoxide or Nitric Anhydride, N 2 5 , is a white 
crystalline solid, which becomes yellow when gently warmed, 
and melts at about 30° to a dark-yellow liquid. It explodes 
when suddenly heated, and sometimes spontaneously at or- 
dinary temperatures. It is prepared by acting on pure nitric 
acid with phosphorus pentoxide : 

2N0 2 -OH + po 2 >0 = ^ 2 >0+ 2P0 2 -OH. 

Nitrogen pentoxide combines with water to form nitric 
acid. It volatilizes in dry air. It combines with nitric acid 

NO.-0-NO-OH 
to form pernitric acid, N„O r .2HNO, or >0 . 

NO-O-NO-OII 

Pernitric acid has a density of 1.64, and crystallizes at 5°. 

Nitrosyl Chloride, N0C1, is formed by the combination o{ 
nitric oxide and chlorine. It is an orange-yellow gas at ordi- 



248 THE FIFTH GROUP. 

nary temperature. It is the chloranhydride of nitrous acid, 
and reacts with a solution of potassium hydroxide, thus : 

N0-C1 + 2K-OH = NO-O-K + K-Cl + H 2 0. 

Nitroxyl Chloride, N0 2 C1, results from the direct union of 
nitrogen tetroxide and chlorine, aided by heat. It is the 
chloranhydride of nitric acid. 

Nitro-Hydrochloric Acid or Aqua Regia is a mixture of 
one measure of nitric acid with two to four measures of hy- 
drochloric acid. The two acids decompose each other; 
chlorine is liberated, and NOC1 and possibly N0 2 C1 are 
formed. Gold and platinum are insoluble in either nitric or 
hydrochloric acid alone, but dissolve in aqua regia. Its 
great solvent power is due to free chlorine, and to a less de- 
gree to NOC1 and N0 2 C1. 

Exp. 171. — Place gold-leaf in fuming hydrochloric acid, and warm. 
Also try to dissolve gold-leaf in hot concentrated nitric acid. Mix the 
two acids containing the gold-leaf. 

Constitution of the Oxygen Compounds of Nitrogen. — Oxy- 
gen acids, as already stated, are regarded as hydroxides, that 
is, compounds of hydroxyl, OH. Assuming that nitric acid^is a 
hydroxide, we have the structural formula IS r 2 -0H. This 
view is supported by the reaction 

N0 2 -C1 + H-OH = NO.-OH + H-Cl, 

in which chlorine is replaced by hydroxyl. 

The radical N0 2 will replace hydrogen not only in water, 
but also in other compounds, as for example : 

C 6 H 6 + N0 2 -0H = C„H -N0 2 + H 2 0. 

Benzene Nitrobenzene 

Eegarding the constitution of the radical NO s little is 



THE ATMOSPHERE. 



known besides its univalent character. 



249 

i 

0=N=0, i^N- 

and 0=N-0- express different views. Nitrons acid is 
NO-OH, and NO is supposed to have the structural for- 
mula = N-. 

The oxides of nitrogen are regarded by some writers as 
having the constitution indicated by the following formulas : 



N=N, 

\/ 


Nitrous oxide 


-N=0, 

Nitric oxide 




0=N-0-N=0, 

Nitrogen trioxide 


_w/0 

iN %0' 


o=n-o-n; 


^0 


^IS-0-N^ 


Nitrogen dioxide 


Nitrogen tetroxide 


Nitrogen pentoxide 



There are also other views of the structure of these com- 
pounds. Nitrogen trioxide and nitrogen pentoxide are the 
anhydrides of nitrous and nitric acids, and may be reasonably 
regarded as containing two acid radicals joined by oxygen. 

Nitrogen- tetroxide dissolves in cold water, with formation 
of nitrous and nitric acids, and hence is to be considered as 
the mixed anhydride of these acids, reacting with water, thus: 

NO-0-N0 2 + HOH = NO-OH + N0 2 -OH. 



The Atmosphere is a mixture of about 21 volumes of oxygen 
and 79 volumes of nitrogen, with a varying proportion of 
aqueous vapor, and an average of 3 volumes of carbon dioxide 
in 10,000. It also contains traces of ammonia, nitrates, and 
nitrites, and at times ozone and hydrogen dioxide. The dust 
of ordinary air, visible when illuminated by a ray of light, is 
partly mineral matter, and partly of animal and vegetable 
origin. With the dust are germs which cause fermentation, 
putrefaction, and at times disease. Liquids, such as mutton- 



250 THE FIFTH GROUP. 

broth and beef -tea, which soon become putrid in ordinary air, 
may be indefinitely preserved by boiling the fluid, and then 
plugging the mouth of the flask with cotton wool so as to 
keep out the dust and germs of the air. The vile odor 
noticed on entering a crowded and poorly-ventilated room is 
due to emanations of putrescible matter from the skin and 
lungs. This is, in part, the cause of the languor and head- 
ache often experienced in an overcrowded room. The in- 
creased amount of carbon dioxide in air which has been 
breathed repeatedly is not harmful, but it indicates the ex- 
tent of the contamination. 

The atmosphere is never free from vapor of Avater. This 
vapor is seldom saturated, that is, in such a condition that any 
slight increase of pressure or diminution of temperature will 
convert a portion of it into water ; but, on the other hand, 
the quantity present is rarely less than one tenth of the 
amount corresponding to saturation. Since the pressure of 
saturated vapor increases rapidly with rising temperature, the 
maximum quantity of water vapor possible in the atmosphere 
increases with the temperature, as is shown by the following 
table : 

Weight of 1 liter of 
saturated water vapor. 

.0126 gram. 
.0487 " 
.0936 « 
.1716 " 
.3010 " 
.5070 " 

By dry air, we understand air which contains but a small 
portion of the water necessary to make a saturated vapor ; and 
by moist or damp air, that which contains a large proportion. 
When the water vapor present is in a condition of saturation, 
the air is commonly said to be saturated. It is obvious that a 



Temperature. 
C. F. 


18° 


0° 


0° 


32° 


10° 


50° 


20° 


68° 


30° 


86° 


40° 


101° 



THE ATMOSPHERE. 251 

dry air at" a high temperature may contain much more water 
than a saturated air at a low temperature ; e.g., a very dry 
air (y 1 ^ saturation) at 40° C. contains four times as much vapor 
as a saturated air at —18° 0. It is for this reason that the 
air in our dwellings is so very dry in winter unless a large 
quantity of water is evaporated in them. For this reason 
also a wetted cloth may lose water very rapidly in an at- 
mosphere which would deposit water upon it at a few 
degrees lower temperature. Air for most experimental pur- 
poses may be dried sufficiently by calcium chloride or concen- 
trated sulphuric acid, but to remove all traces of water long 
contact with phosphorus pentoxide is necessary. 

The weight of one liter of pure dry air at 0° and 760 mm. 
is 1.2936 grams. The density of gases is very commonly re- 
ferred to air as unity. In this book, as already stated, the 
density of gases is referred to hydrogen gas as unity. Dry air 
is 14.44 times heavier than hydrogen, hence the density, 
air = 1, multiplied by 14.44, equals the density compared 
with hydrogen. 

The nitrogen and oxygen of the atmosphere are not in 
chemical combination, as is evident from the following facts. 
The two elements are not in atomic proportions, and they are 
not found to be in exactly the same proportion at all times 
and in all localities. When nitrogen and oxygen are mixed, 
there is no change in temperature or other evidence of chemi- 
cal union. Pure water in contact with air absorbs oxygen 
and nitrogen in other proportions than found in the atmos- 
phere, which would not be the case were the two gases in 
combination. In chemical deportment air is a mixture 1 , its 
oxygen entering into combination the same as pure oxygen. 
Nitric oxide forms when mixed with air red fumes of the 
dioxide, but if mixed with nitrous oxide gas no red fumes ap- 
pear, that is, nitric oxide combines with free oxygen, but does 
not withdraw oxygen from the compound N.,0. The oxygon 
of the air is the supporter of respiration ami ordinary com- 



252 THE FIFTH GEOUP. 

bustion, the chief products in both cases being carbon dioxide 
and water. 



Phosphorus, P. 

Atomic Weight, 31. Molecule, P 4 . 

Phosphorus is a widely distributed element, occurring only 
in combination as in phosphates. It is an essential constitu- 
ent of fruitful soils, and plants will not grow in soil free from 
it. From plants phosphorus passes to animals, whose juices 
and tissues contain small quantities of it, and whose bones are 
largely composed of calcium phosphate. 

Phosphorus was first isolated by Brand in the year 1669, 
who obtained it from urkie. Urine contains phosphates, and 
when evaporated a residue remains, from which phosphorus 
may be distilled out of contact with air. This was the only 
method of preparing phosphorus up to 1771, when Scheele 
obtained it from bones. Bones have since been the chief 
source of the phosphorus of commerce. In the manufacture 
of phosphorus bones are burned to remove the organic matter, 
and there remains bone-ash, which is about four fifths calcium 
phosphate, the rest being mostly calcium carbonate. The 
bone-ash is treated with sulphuric acid, when an insoluble 
calcium sulphate and a soluble acid calcium phosphate are 
formed, thus: 

Ca 3 (P0 4 ) 2 + 2H 2 S0 4 = CaH 4 (P0 4 ) 2 + 2CaS0 4 . 

The solution of the acid calcium phosphate is filtered, evapo- 
rated to dryness, and the residue is heated to change the acid 
salt into the metaphosphate: 

CaH,(P0 4 ), = Ca(PO,) a + 2H a O. 



PHOSPHORUS. 253 

The calcium metaphosphate is then intimately mixed with 
charcoal and heated to redness in earthen retorts, and the 
phosphorus vapor is condensed in water: 

3Ca(P0 3 ) a + 100 = P 4 + Ca 3 (P0 4 ) 2 + 1000. 

To obtain all of the phosphorus, sand is added to the mixture, 
and a calcium silicate remains in the retort. Phosphorus is 
also manufactured from mineral phosphates. 

There are three allotropic forms of solid phosphorus, the 
common or octahedral, the red or amorphous, and the metallic 
or rhombohedral. The gas density of phosphorus at 1040° 
has been found to be 65; hence the molecule of phosphorus 
contains four atoms. At a white heat (1437°) Mensching and 
Meyer found the gas density to be 43. 8. This result shows 
that the molecules of P 4 dissociate into molecules containing 
fewer atoms. 

Common or Octahedral Phosphorus is almost colorless, and 
when slowly cooled from the liquid state is transparent, but if 
quickly cooled it is wax-like and translucent. At ordinary 
temperature it may be cut with a knife, but when cooled it is 
brittle. It has a density at 10° of 1.83. It melts at 44.3° 
and boils at 278°, forming a colorless vapor. The best solvent 
for common phosphorus is carbon disulphide. It is slightly 
soluble in some oils, but is insoluble in water and alcohol. 
Phosphorus ignites in air when heated to 60°, and may be in- 
flamed by friction or a warm rod. Because of its inflamma- 
bility it is kept under water. The familiar luminosity in the 
dark, or phosphorescence of phosphorus, from which it derives 
its name, is accompanied by a slow oxidation, and is host 
seen in a warm, damp atmosphere. A luminous mist rises 
from the phosphorus, and a peculiar odor is perceptible. It 
is remarkable that phosphorus is not luminous in pure oxygen 
Ibelow 20°, but on diluting with an indifferent gas or reducing 



254 THE FIFTH GROUP. 

the pressure of the oxygen to \ of an atmosphere the phenom- 
enon of phosphorescence is observed. According to Baker, 
phosphorus may be melted, and even distilled without taking 
fire in oxygen, which has been dried by long contact with 
phosphorus pentoxide to free the gas from all traces of mois- 
ture. From this it appears that water, even when present in 
minute quantity, plays some part in the combustion of phos- 
phorus in oxygen or air. 

Sticks of phosphorus which are kept under water and ex- 
posed to light become covered, first with an opaque white 
coating which gradually turns red and flakes off. This is red 
phosphorus. A solution of phosphorus in carbon disulphide 
on standing in sunlight deposits 3^ellow particles of amor- 
phous phosphorus, which later turn reddish. 

Common phosphorus is a poison, large doses causing death 
in a few hours, and small doses a lingering and painful ill- 
ness. The fumes are poisonous, and persons constantly ex- 
posed to them are liable to phosphorus-necrosis, a disease in 
which the bones of the jaw waste away. Red phosphorus is not 
poisonous. Phosphorus burns are painful, and often quite 
severe. They should be soaked for half an hour or longer in 
lime water to neutralize the acid formed in the combustion of 
the phosphorus, and then treated like an ordinary wound. 

Red or Amorphous Phosphorus is formed when common 
phosphorus is heated to 240°-250°. Above 260° the red phos- 
phorus changes to the colorless modification. The presence 
of a small quantity of iodine greatly facilitates the formation 
of red phosphorus. The change takes place rapidly when 
colorless phosphorus is heated in a closed vessel to about 300°, 
or above its boiling point. Thus obtained, red phosphorus is 
mixed with common phosphorus, which is removed by extrac- 
tion with carbon disulphide. 

The color of amorphous phosphorus is brownish red, but 
when finely divided and suspended in water it appears of a 



PHOSPHORUS. 255 

bright scarlet color. Eed phosphorus has a density of 2.1, is 
insoluble, and non-volatile. It is not phosphorescent, and 
does not oxidize in air. In its chemical deportment it is less 
energetic than ordinary phosphorus. 

Rhomb ohedral Phosphorus is obtained by heating phos- 
phorus with lead, in a sealed tube, to a red heat. After cool- 
ing, the lead is dissolved in nitric acid, and a mass of dark- 
red crystals remains. This modification of phosphorus has a 
density of 2.34, and requires a temperature of 358° to convert 
it into the ordinary form. 

Phosphorus is used in chemical industries and in the labora- 
tory, but its chief use is in the manufacture of matches. 
Friction matches were invented in 1832. Before this time 
fire was obtained by igniting tinder with sparks from a flint 
and steel, and also by igniting a mixture of potassium chlo- 
rate and sugar, to which sulphur was sometimes added, with 
strong sulphuric acid. The sulphur match of the present 
time is made" by dipping the match-stick into sulphur, and 
then tipping the end with a mixture of phosphorus, glue, 
potassium nitrate or chlorate, and other substances. The mix- 
ture inflames when rubbed, and ignites the sulphur, which in 
turn sets fire to the wood. Paraffin is now substituted for sul- 
phur to avoid the disagreeable sulphur fumes. 

Exp. 172. — To purify common phosphorus place it in a dilute solution 
of potassium dichromate and sulphuric acid, and warm on a water-bath 
for some time. The impurities will be oxidized, and most of the coat- 
ing which was on the phosphorus will disappear. Wash the solution 
from the beaker by means of a stream of water, and if the phosphorus 
is not colorless and clean, repeat the treatment. Melt it again under a 
fresh solution of dichromate and acid, and allow to cool slowly and 
quietly. Frequently it will remain liquid for a, long time at a tempera- 
ture much below its melting point. Phosphorus thus treated is nearly 
white, translucent, and exhibits on its upper surface a fern-like crystal- 
line structure. 



256 THE FIFTH GROUP. 

Exp. 173. — Melt some clean phosphorus under water in a beaker, and 
draw it (best by means of a rubber bulb) into a glass tube which is 
slightly conical. Pour cold water on the tube to solidify the phos- 
phorus, and then transfer the tube to a beaker of cold water, and press 
out the stick of phosphorus. Preserve it in a bottle of water, and ob- 
serve that after a day or two its surface becomes opaque. 

Exp. 174. — Dissolve a clean piece of phosphorus in carbon disulphide 
and pour some of the solution upon filter paper. The solvent will 
soon evaporate, and the paper will glow in the dark and soon inflame. 
The solution of phosphorus will become turbid on standing in daylight. 

Exp. 175. — Dry some phosphorus with filter paper, and heat it in a 
dry test-tube to boiling. But little red phosphorus will be formed. Add 
a fragment of iodine and heat again. 

Exp. 176.— Examine commercial red phosphorus, and notice that it is 
moist, and that it reddens blue litmus paper. It contains a little com- 
mon phosphorus, which oxidizes and absorbs moisture from the air, 
forming acids. 

Heat a little red phosphorus in a narrow glass tube closed at one end, 
to convert it into the common modification. At first a gas may escape 
and take fire spontaneously. It is phosphoretted hydrogen, formed 
from the decomposition of phosphorous acid contained in the moist 
red phosphorus. 

There are three compounds of phosphorus and hydrogen, 
PH 3 , P 2 H 4 , and P 4 H 2 . The first is the analogue of ammonia, 
but is less basic. The second corresponds to hydrazine, N 2 H 4 . 

Phosphine or Phosphoretted Hydrogen, PH 3 , is a colorless 
gas with a peculiar disagreeable odor, and is but slightly 
soluble in water. It condenses to a liquid, boiling at about 
— 85°, and solidifying at —133°. When pure it is not spon- 
taneously inflammable. It is formed in a number of reactions, 
and is best obtained pure by decomposing phosphonium iodide 
with a solution of potassium hydroxide: 

PHJ + KOH = PH 3 + KI -f H 2 0. 

The gas density of phosphine corresponds to the formula 



LIQUID HYDROGEN" PHOSPHIDE. 



257 



PH 3 . Phosphine is readily decomposed by electric sparks, 
and two volumes of the gas yield three volumes of hydrogen 
and half a volume of phosphorus vapor. Phosphine does not 
neutralize acid solutions to form salts, but unites with hydri- 
odic acid gas, forming phosphonium iodide, PH 4 I, in colorless 
crystals. This compound is decomposed by water into PH 3 
and HI. 

Liquid Hydrogen Phosphide, P 2 H 4 , is an unstable compound 
which takes fire in air. It is formed in small quantity, together 
with phosphine, as in the following experiment. 

Exp. 177. — Place in a flask, Fig. 84, a solution of potassium hydrox- 
ide, density 1.30, a piece of phosphorus and a few drops of ether. The 
latter serves to displace the air in the flask. Heat cautiously until gas 
escapes freely from the delivery-tube, which dips under water. In 
order to obtain large bubbles of gas, a thin wide tube, made from a test- 




Fig. 84. 



tube, is fastened by means of a flexible rubber tube to the delivery- tube. 
Many of the bubbles will inflame spontaneously, and form beautiful 
smoke-rings. The water in the pan should not be below ordinary tern- 



258 THE FIFTH GKOUP. 

perature, since cold water may condense the liquid hydrogen phos- 
phide, to which is due the spontaneous ignition of the gas. 

The gas formed in the experiment is a mixture of phosphine, hydro- 
gen, and liquid hydrogen phosphide. 



Phosphorus Halides. 

Phosphorus Trifluoride, PF 3 , is a gas obtained by the action of copper 
phosphide on lead fluoride. It condenses at low temperature and high 
pressure to a colorless liquid. It is decomposed slowly by water, with 
formation of phosphorous and hydrofluoric acids. A mixture of 2 
volumes of the gas and 1 volume of oxygen explode violently when 
ignited by the electric spark, and form 2 volumes of the oxyfiuoride, 
POF 3 . 

Phosphorus Pentafluoride, PF 6 . — This compound is of great 
theoretical interest, as it shows the pentavalent character of 
phosphorus. It is formed by the action of arsenic trifluoride 
on phosphorus pentachloride : 

5AsF 3 + 3PC1 5 = 5AsCl 3 + 3PF 5 . 

It is a colorless gas, possessing a density of 63. The molecular 
weight is, therefore, 126, corresponding to PF 5 . 

Phosphorus Trichloride, PC1 3 , is prepared by acting on an 
excess of phosphorus with dry chlorine. It is a colorless, 
mobile liquid, boiliug at 76°, and having a gas density of 68.5. 
With water it forms phosphorous and hydrochloric acids. 

Phosphorus Pentachloride, PC1 5 , results from the union of 
chlorine with the trichloride. It is a yellowish crystalline 
powder. It dissociates gradually when vaporized into equal 
molecules of PC1 3 and Cl 2 . In the presence of the vapor of 
the trichloride the pentachloride is more stable, and its vapor 
has been found to have a density nearly that required by the 
formula, PC1 6 . Phosphorus pentachloride changes in moist 



PHOSPHORUS HALIDES. 259 

air to the oxychloride, and with excess of water or triphos- 
phoric acid is formed. 

Phosphorus Oxychloride, Phosphoryl Chloride, P0C1 3 or 

/ cl . 

P0^-C1. — This compound is a colorless fuming liquid, boiling 

at 107°, and with a gas density corresponding to its formula. 
It dissolves slowly in water, with formation of orthophosphoric 
and hydrochloric acids. It is formed in a number of reac- 
tions, of which may be mentioned the following : 

PC1 5 + H 2 = POC1 3 + 2HC1. 
3PC1 5 + P 2 5 = 5P0C1 3 . 

The bromides of phosphorus are analogous to the chlorides. 

Phosphorus Di-iodide, PI 2 or P 2 I 4 , is prepared by dissolving 
phosphorus and iodine in proper proportions in carbon disul- 
phide. On cooling the solution yellow crystals of the di-iodide 
separate, or the carbon disulphide maybe evaporated, and the 
last traces of it expelled by a current of dry air. It is decom- 
posed by water, with formation of phosphorous and hydriodic 
acids and a yellow substance. 

Phosphorus Tri-iodide, PI 3 , is a red compound obtained in 
the same way as the foregoing, only using more iodine. It 
reacts with water, thus : 

PI, + 3H a O = P(OH), + 3HI. 

The phosphorus halides are valuable reagents for making 
other compounds, and are especially useful in effecting the re- 
placement of hydroxy!, OH, by a halogen atom. Examples of 
the reaction will be given under the constitution of phosphoric 
acids. There is a gradation of properties in these halides. 
with increasing atomic weights of the halogens. The penta- 
fluoride is quite stable, the pentaehloride dissociates on heat- 



2G0 



THE FIFTH GROUP. 



iiig, and the pentabroinide more easily, while the pentiodide 
of phosphorus does not exist. 



Oxides and Hydroxides of Phosphorus. 



Phosphorus suboxide, P 4 0. 
Phosphorous anhydride, P 2 3 . 

Phosphorus tetroxide, P 2 4 . 



Phosphoric anhydride, P 2 O f 



Hypophosphorous acid, H 2 PO-OH 

/OH 

Phosphorous acid, P^-OH 

\OH 



Hypophosphoric acid, 

Orthophosphoric acid, 

Pyrophosphoric acid, 
L Metaphosphoric acid, 



,PO< 
P < 



OH 
OH 
OH 
OH 



/OH 

POf OH 

\OH 



,PO< 
PO< 



OH 
OH 
OH 
OH 



P0 2 -OH 



The oxides of phosphorus, with the exception of P 2 O s , are 
little understood, and are obtained in the pure state with diffi- 
culty. These lower oxides are formed when phosphorus is 
burned with an incomplete supply of air. 

Phosphorus Pentoxide or Phosphoric Anhydride, P 2 0., is 

prepared by burning phosphorus in dry air. It is a white 
amorphous powder which rapidly absorbs moisture from the 
air and deliquesces. It hisses when thrown upon water and 
dissolves with formation of metaphosphoric acid: 

P.O. 4- HO = 2HPO„. 



Phosphorus pentoxide is used for drying gases, and for re- 
moving the elements of water from acids in the preparation 






PHOSPHOROUS ACID. 261 

of acid anhydrides, as, for example, N 2 5 . It is a valuable 
reagent in organic chemistry. 

Exp. 178. — Ignite a piece of dry phosphorus in a capsule standing on 
a plate, and cover with a bell-jar. After the snow-like powder has 
settled, remove the bell- jar and capsule, and pour a little water on the 
phosphorus pentoxide which has collected on the plate. Test the solu- 
tion with blue litmus paper. The white powder adhering to the sides 
of the bell- jar will soon liquefy. 

Hypophosphorous Acid, H 3 P0 2 or H 2 P0-0H. — The barium 
salt of this acid is prepared by heating phosphorus with a 
solution of barium hydroxide as long as phosi3hine escapes: 

3Ba(OH) 2 + 8P + 6H 2 = 3Ba(H 2 P0 2 ) 2 + 2PH 3 . 

The excess of barium hydroxide is converted by a stream of 
carbon dioxide into carbonate, which is removed by filtration. 
The solution yields on evaporation crystals of Ba(P0 2 H 2 ) 2 -j- 
H 2 0. The acid is obtained as follows. Sulphuric acid is 
added to a solution of the barium salt, and the mixture 
allowed to stand until the precipitated barium sulphate has 
subsided. The solution is siphoned off and evaporated, the 
temperature being raised at last to 130°. Crystals separate on 
cooling the liquid in a freezing mixture. The pure acid melts 
at 17°. 4. Hypophosphorous acid and solutions of its salts ab- 
sorb oxygen, and change to phosphoric acid and phosphates. 
The acid precipitates the metals gold, silver, and mercury 
from solutions of their salts. But one atom of hydrogen in 
hypophosphorous acid can be replaced by basic radicals, and 
therefore the acid must be considered monobasic. 

Phosphorous Acid, H,P0 3 , HP0(GII) 2 or P(0H) 3 .— This 

acid is prepared by adding gradually phosphorus trichloride 
to water: 

PCI., + 3H a O = H 8 PO a -f 3IIC1. 

The vessel is placed in ice water to prevent the solution be- 



262 THE FIFTH GROUP. 

coming too hot from the reaction. The solution is evaporated 
and finally heated to 180°. The thick liquid obtained solidi- 
fies after a time to a crystalline mass. Phosphorous acid melts 
at 70° and deliquesces in air. When heated it decomposes 
thus: 

4H 3 P0 3 = 3H 3 P0 4 + PH 3 . 

Its aqueous solution absorbs oxygen from the air, and re- 
sembles hypophosphorous acid in reducing action. It is 
dibasic under ordinary circumstances, forming salts, such as 
Na 2 HP0 3 . There are, however, organic derivatives of a tri- 
basic phosphorous acid, as for example tri-ethyl phosphite, 
P(OC 2 H 5 ),. 

Hypophosphoric Acid, H 4 P 2 6 , is formed together with phosphorous 
and phosphoric acids by the slow oxidation of phosphorus in moist air. 
The sodium salts are Na 4 P 2 8 -f- 10H 2 O, Na 3 HP 2 6 + 9H 2 0, and 
Na 2 H 2 P 2 6 + 6H 2 0. 

Phosphoric Acids. — Phosphorus pentahydroxide, P(OH) 5 , 
the hypothetical hydroxide of pentavalent phosphorus, is un- 
known; but compounds, which may be viewed as derived 
from it by removal of the elements of one or more molecules 
of water, are well-known acids, viz. : 

O^ Vn / ATT 

P 1 OH - HOH = P^CCxS 

OH U 



OH 
-OH 



Orthophosphoric acid 

/-OH 
- HOH = P^^O 



Metaphosphoric acid 



PHOSPHORIC ACIDS. 263 



/ 0H cm 

- HOH = 

"OH 

Diphosphoric acid 



OR 
.OH 
\OH 



o=Pf-OH o=]p<° 




Ortliopliosplioric acid is tribasic, and forms salts in which all 
three hydrogen atoms are replaced, and it is therefore assumed 
to contain three hydroxyl groups. This accords with the sim- 
plest view of the reaction between the chloranhydride, P0C1 3 
and water, thus: 

" OH .OH 

OH = POf-OH + 3HC1. 
OH \OH 

The formula PO(OH) 3 represents all that is known of the 
constitution of orthophosphoric acid, and does not require us 
to assign a valence to phosphorus, and thus a structure to the 
radical PO. 

Metaphosphoric acid may be viewed thus, i ^P-OH, 

or the valence of phosphorus may be left undecided, and the 
structural formula P0 2 -OH given. It will be observed that 
this is analogous to nitric acid. Diphosphoric acid, because of 
its derivation from the ortho-acid, may be assumed to contain 
the radical PO, and, without expressing the valence of phos- 
phorus, it may be formulated thus: 

PO< oii 

0/ OH 

iu< 011 

It is remarkable that either two or four, and never three. 



264 THE FIFTH GROUP. 

atoms of hydrogen are replaced in the formation of diphos- 
phates. There are salts which correspond to more complex 
phosphoric acids. 

The orthophosphates are distinguished by the following re- 
actions. Silver nitrate produces in solutions of the normal 
and acid salts a yellow precipitate of silver orthophosphate, 
Ag 3 P0 4 . Ammonium molybdate throws down from a solution 
of orthophosphate in nitric acid a heavy yellow precipitate of 
ammonium phospho-molybdate. 

Metaphosphoric acid coagulates a solution of albumen, the 
other phosphoric acids do not. 

Diphosphoric acid is distinguished from orthophosphoric by 
producing a white granular precipitate in a solution of silver 
nitrate, and from metaphosphoric acid by the fact that it does 
not precipitate albumen. 

/OH 

Orthophosphoric Acid, H 3 P0 4 or PO^-OH, is obtained by 

M)H 

dissolving common phosphorus in warm nitric acid of 
density 1.20. A more concentrated nitric acid should not 
be used on account of danger of explosion. The solution is 
evaporated and heated as long as acid fumes escape. Phos- 
phoric acid is manufactured from bone-ash, the calcium 
being separated as sulphate by means of sulphuric acid. The 
process yields a product containing a small quantity of mag- 
nesia, and often other impurities. When an aqueous solution 
of phosphoric acid is evaporated, and the residue heated to 
150°, but not hotter, a viscous mass remains, which has the 
composition H 3 P0 4 , and which crystallizes on standing, or 
quickly if a crystal of the acid is added. The crystals melt at 
38°. 6, and deliquesce in moist air. A solution of phosphoric 
acid has an agreeable acid taste. 



PHOSPHATES. 265 

/O-Na 
Trisodium Orthophosphate, Na 3 P0 4 or P0~0-Na. — This salt 

\0-Na 

is prepared by adding the required amount of sodium 
hydroxide to a solution of disodium phosphate. On evapora- 
tion crystals of Na 3 P0 4 -j- 12H 2 separate. These lose water 
on heating, and are converted into the anhydrous salt, which 
is not decomposed at a red heat. A solution of trisodium 
phosphate reacts alkaline, and absorbs carbon dioxide from the 
air. The reaction in presence of sufficient carbon dioxide is 
as follows: 



Na 3 P0 4 + C0 2 + H 2 = HNa 2 P0 4 + HNaCO,. 

/O-Na 
Disodium Orthophosphate, Na 2 HP0 4 or P0(-0-Na, is the com- 

M)H 

mon phosphate of soda. It is prepared by adding sodium 
carbonate to a solution of orthophosphoric acid as long as car- 
bon dioxide escapes. The solution yields on evaporation crys- 
tals having the composition, Na 2 HP0 4 + 12H 2 0. The salt is 
very soluble in water, and, although an acid salt, reacts alka- 
line to litmus. It is much used in the preparation of other 
phosphates. 

/O-Na 
Monosodium Orthophosphate, NaH 2 P0 4 or PO^OH . — This 

M)H 

compound results when solutions of common sodium phos- 
phate and orthophosphoric acid are mixed. Crystals contain- 
ing one molecule of water of crystallization separate when the 
solution is evaporated. The salt has an acid reaction. 

Silver Orthophosphate, Ag.,P0 4 , separates as a yellow pre- 
cipitate on adding silver nitrate to solutions of either of the 
sodium orthophosp hates : 



266 THE FIFTH GEOUP. 

Na,P0 4 + 3AgNO, = Ag,P0 4 + 31Ms T 3 . 
Na a HP0 4 + 3AgNO, = Ag,P0 4 + SHSaNO, + HM) 3 . 
NaH a P0 4 + 3AgN0 3 = Ag 3 P0 4 + NaNO s + 2HNO,. 

The nitric acid set free in the last two reactions renders the 
solutions acid. Silver phosphate is soluble in nitric acid and 
in ammonia. 

Ammonium Sodium Orthophosphate, NH 4 NaHP0 4 or 

/0-NH 4 
PO^O-Na . — This compound is commonly known as micro- 
NDH 

cosmic salt. It is obtained by mixing solutions of sodium 
and ammonium phosphates, and also by the following process. 
Common sodium phosphate, 6 or 7 parts, and 2 parts of sal- 
ammoniac, are dissolved in boiling water. The salt separates 
in crystals with four molecules of water as the solution cools. 
The mother-liquor contains common salt. 

Na 2 HP0 4 + NH 4 C1 = NH 4 ffaHP0 4 + NaCl. 

By recrystallization with addition of ammonia the salt is 
freed from sodium chloride. It loses its water at a gentle 
heat, and at a red heat fuses and is converted into sodium 
metaphosphate. Microcosmic salt is used as a flux in blow 
piping. 

/0 >Ca 

P0 \0 

Tricalcium Orthophosphate, Ca 3 (P0 4 ) 2 or > n >Ca. — This 

\g>Ca 

calcium phosphate separates as a gelatinous hydrous precipi- 
tate on addition of sodium phosphate to an ammoniacal solu- 
tion of calcium chloride. It is the chief mineral constituent 
of bones. It occurs abundantly in nature in combination 
with calcium chloride and fluoride in the mineral apatite. 



PHOSPHATES. 267 

It is insoluble in water, and is decomposed by sulphuric acid 
with formation of monocalcium phosphate and calcium sul- 
phate. 

Dicalcium Orthophosphate, Ca 2 H 2 (P0 4 ) 2 or Hydrogen Cal- 
cium Orthophosphate, HCaP0 4 . — This compound is precipi- 
tated when calcium chloride is added to a solution of sodium 
phosphate. 

Tetrahydrogen Calcium Orthophosphate, H 4 Ca(P0 4 ) 2 or 
/OH 
PO^-OH 

Q>Ca. — This salt is obtained by dissolving either of the 

PO^-OH 
M)H 

foregoing calcium phosphates in the requisite amount of 
phosphoric acid. It crystallizes, on spontaneous evaporation, 
in scales containing one molecule of water. It is also formed 
when tricalcium phosphate is treated with acids. The fertil- 
izer known as " superphosphate of lime" is made by treating 
bones with sulphuric acid It is a mixture of tetrahydrogen 
calcium phosphate, calcium sulphate, and organic matter. 
Tetrahydrogen calcium phosphate dissolves without decom- 
position in much water, but when 10 to 40 parts of water are 
added, it decomposes with separation of hydrogen calcium 
phosphate. Also, on heating the concentrated aqueous solu- 
tion the hydrogen calcium phosphate separates, and redissolves 
again on cooling. 

Ammonium Magnesium Orthophosphate, NH.MgPO, -J- 

/0-NH 4 
6H 2 or P0~0^ Mo , + 6H o 0.— This double salt separates as a 
\q > 1V1 & 

white precipitate when an ammoniacal solution of magnesium 
chloride and ammonium chloride is added to a solution of an 

orthophosphate. The precipitate is insoluble in water con- 



268 THE FIFTH GKOUP. 

taming ammonia, but is readily soluble in acids. It is ol> 
tained in the separation of magnesium, and also of phosphorus, 
in chemical analysis. It is converted by ignition into mag- 
nesium pyrophosphate, Mg^O^ 

P0< 0H 
/OH 

Diphosphoric, Pyrophosphoric Acid, H 4 P 2 7 or 0< nTT . 

\PO<^ 

— This acid is formed when orthophosphoric acid is heated for 
some time to 215°. The sodium salt is easier to prepare pure, 
and is made by heating common sodium phosphate : 

/O-JSTa 
PO^-O-Xa p n ^O-]S3"a 

OH - M > U - U \ p() O-Na' 

\0-Xa 

Sodium pyrophosphate is soluble in water, and does not 
change to orthophosphate on boiling unless a stronger acid is 
added. When silver nitrate is added to a solution of sodium 
pyrophosphate, a white precipitate of silver pyrophosphate, 
Ag 4 P 2 T , separates. 

Metaphosphoric, Glacial Phosphoric Acid, HP0 3 or P0,-0H, 

remains as a glassy mass when pyrophosphoric acid is heated to 
redness. The molten mass solidifies on cooling to a colorless 
glass. Metaphosphoric acid dissolves slowly in water, and 
deliquesces in moist air. In solution it changes slowly, or 
rapidly if heated, to orthophosphoric acid. 



Arsenic * As. 

Atomic Weight, 75. Molecule, As 4 . 

Arsenic occurs in the free state in abundance in some 
localities, but is more commonly found in combination, as in 



ARSENIC. 269 

the minerals mispickel, FeSAs, realgar, As 2 S 2 , and orpiment, 
As a S 3 . Arsenic is found in most iron pyrites, and is often 
contained in sulphuric acid manufactured from pyrites, and 
in substances prepared by means of such acid, 

The metallic arsenic of commerce is either the native ar- 
senic from the mines, or the product obtained by heating 
mispickel, which is decomposed thus : 

FeSAs = FeS + As. 

It usually has a dark-gray coating, due to oxidation. It may 
be purified by sublimation out of contact with air. Metallic 
arsenic is sold under the name of cobalt, a term derived from 
the German word Kohold, the demon of the mines. Cobalt, 
the element resembling nickel, is usually found in combina- 
tion with arsenic. 

Arsenic has a grayish-white metallic lustre, a density of 
5.7, and is crystalline and very brittle. It volatilizes when 
heated without previous fusion, but may be fused at a dull- 
red heat, under increased pressure, in a sealed tube. The 
vapor of arsenic has a yellow color and a disagreeable garlic 
odor, supposed by some writers to be due to a lower oxide of 
arsenic. The gas density of arsenic at a red heat has been 
found to be 150, corresponding to the molecule As 4 . The 
density diminishes with increasing temperature, and at a 
white heat (1437°) Mensching and Meyer have found it to be 
94.5. 

There are two, and perhaps three, amorphous modifications 
of arsenic. When arsenic is sublimed in a current of hydro- 
gen, or in a tube sealed at one end, it condenses near the 
heated part in the crystalline form, and beyond as a black 
vitreous coating, while further on a yellow deposit* is formed, 

* The writer has observed that commercial arsenic gives a yellow 
sublimate on heating, but that a sample, which has been heated until it 



270 THE FIFTH GKOUP. 

according to Bellendorff, which quickly turns gray. These 
two forms of amorphous arsenic have a density of 4.7, and 
change at about 360° into the crystalline form. 

Arsenic oxidizes slowly in moist air, losing its metallic 
lustre, and becoming dark-gray. According to Greuther, the 
product of the oxidation is arsenious oxide. Heated in air or 
oxygen, arsenic burns with a peculiar white flame, arsenious 
oxide being formed. It dissolves readily in nitric acid. 

Exp. 179. — Heat a little arsenic in a tube closed at one end until it is 
bright, and after it is cool transfer to another tube. Sublime it slowly, 
and observe the three modifications. If the black and gray amorphous 
forms are gently heated they will become crystalline. 

Arseniuretted Hydrogen, Arsine, AsH 3 , is a colorless, highly 
poisonous gas, with an odor somewhat different from that of 
arsenical vapor. It is obtained by dissolving zinc arsenide, 
Zn 2 As 3 , in dilute sulphuric acid, or in greater purity when 
sodium arsenide, Na 3 As, is dissolved in water or very dilute 
hydrochloric acid. 

Arsene is formed in acid solutions containing arsenic, from 
which hydrogen is evolved, and passes off with the hydrogen. 
Arsenious oxide and arsenic acid are reduced by nascent hy- 
drogen, as follows : 

As 4 6 + 24H = 4AsH 3 + 6H 2 0. 
H,As0 4 + 8H = AsH 3 + 4H 2 0. 

Arsine burns with a bluish-white flame, with formation of 
water and arsenious oxide. It is decomposed by heat into its 
constituents. Chemically, it closely resembles phosphine, but 
differs from the latter in that it does not combine with acids 
to form salts. 

remains bright on cooling, does not give a yellow -colored deposit. 
Whether there is a yellow modification of arsenic, or whether the sub- 
limate in question is due to impurities, requires investigation, 



AESENIC. 



271 



Exp. 180. — Place 20-30 grams of arsenic-free granulated zinc in the 
generator of the Marsh apparatus, Fig. 85, and pour in dilute sulphuric 
acid, and a drop or two of platinum solution to facilitate the action be- 
tween the acid and zinc. The escaping gas is dried by means of the 
calcium chloride tube. The portion of the hard glass tube over the 
lamps is wrapped with wire gauze to keep it in shape when hot. After 
the air is completely expelled from the apparatus ignite the jet of hydro- 
gen, and heat the tube by the lamp flames. If no dark deposit appear 




Fig. 85. 



in the tube beyond the heated portion after some time, the reagents are 
free from arsenic. Next add a solution containing a few milligrams 
of arsenious oxide to the acid in the funnel tube, and allow this acid 
solution to drop slowly into the generator. After a time an arsenical 
mirror will be seen in the tube. This is amorphous arsenic. If the 
current of hydrogen is not too rapid, and but a few milligrams of arsenic 
are added to the acid in the generator, all of the metal will be de- 
posited in the tube in from two to four hours. To find the weighl of 
the arsenic cut off the length of tube containing the mirror ami weigh. 
The metal can then be removed by heat, or by solution in nitric acid, 
and the tube cleaned, dried, and again weighed. The difference iu the 
two weights gives the amount of metallic arsenic. This method of de- 
tecting arsenic is remarkable for its extreme delicacy, , „'„„ of one milli- 
gram of arsenious^oxide producing in this form of the Marsh apparatus 
a distinct mirror. 

After obtaining a deposit in the tube from a few milligrams of ar- 



272 THE FIFTH GROUP. 

senic, add a solution of 20-30 milligrams of arsenious oxide to the 
generator, and sufficient acid to make the hydrogen flame half an inch 
or longer. The flame will have a whitish tinge, and if a cold piece of 
porcelain is held in the flame a dark coating of arsenic will be deposited. 
The AsH 3 is decomposed by the heat of the flame, and part of the metal 
condenses on the cold surface. 

Arsenic Trifluoride, AsF 3 , is prepared by distilling in a leaden vessel 
a mixture of equal parts of calcium fluoride and arsenious oxide with 
four parts of pure sulphuric acid free from water. It is a colorless, 
fuming liquid, boiling at 63°. In contact with the skin it produces 
deep, painful sores. 

Arsenic Trichloride, AsCl 3 , is formed by the direct union of its ele- 
ments. It is a very poisonous, fuming liquid, with a boiling point of 
134 c . Its observed gas density corresponds to the formula AsCl 3 . It 
mixes with a little water, and is decomposed by much water, with for- 
mation of hydrochloric acid and arsenious oxide, which separates in 
minute crystals. A solution of arsenious oxide in concentrated hydro- 
chloric acid probably contains arsenic trichloride. 

Arsenic Di-iodide, Asl 2 , is a dark-red substance which is soluble in 
alcohol and ether. It oxidizes readily on exposure to air, and is decom- 
posed by water, with separation of arsenic in the form of a black 
powder. Arsenic Tri-iodide, Asl 3 , is prepared by subliming a mixture 
of 1 part of metallic arsenic and 3 parts of iodine. It forms brilliant red 
crystals, soluble in water and alcohol. It is used in medicine. Arsenic 
Pentiodide, Asl 5 , is obtained, according to Sloan, by heating arsenic and 
iodine in the proportions required by the formula to 150° for an hour and 
a half in a vessel filled* with carbon dioxide. 



Oxides and Hydroxides of Arsenic. 

Arsenious anhydride, As 4 6 . /OH 

Orthoarsenic acid, AsO^- OH 

\OH 

/AsO<Q H 
Arsenic anhydride, A 2 5 . Diarsenic acid, 0<^ qtt 

\AsO< OH 

Metarsenic acid, AsO^-OH. 



ARSENIOUS OXIDE. 273 

The aqueous solution of arsenious oxide doubtless contains 
arsenious acid, which has not, however, been isolated. There 
are salts of orthoarsemous acid, As(OH) 3 , and of metarsenious 
acid, AsO.OH. 

Arsenious Oxide or Anhydride, As 4 6 , is the arsenic or 
white arsenic of commerce. It is also called arsenic trioxide 
and arsenious acid. It is obtained in large quantities, as a 
waste product, in the roasting of arsenical ores, the vapors of 
it being condensed in large chambers. The crude product is 
purified by sublimation. If the vapors are conducted into a 
receiver, where the temperature is but little beiow that re- 
quired to volatilize the oxide, vitreous arsenic, the amorphous 
form, is obtained ; but when the condensation occurs in a cool 
chamber, a white crystalline powder results. Both forms of 
arsenious oxide are commercial articles. 

The amorphous form of arsenious oxide is a transparent, 
colorless solid, like glass in appearance. It gradually becomes 
opaque white, owing to the change to the crystalline condition. 
There are t\vo crystalline forms of arsenious oxide, the rhombic, 
which is occasionally obtained by sublimation, and the octahe- 
dral, which commonly results when the vapor is cooled. Arse- 
nious oxide volatilizes at about 218°, forming a colorless vapor. 
Its gas density has been determined at temperatures from 571° 
to 1560°, and found to bo nearly 198, the calculated gas density 
of the compound As 4 6 . Arsenious oxide is but slightly 
soluble in water, the vitreous variety being more soluble than 
the crystalline. It dissolves abundantly in hot concentrated 
hydrochloric acid, and the solution on cooling deposits octa- 
hedral crystals. It forms with fuming sulphuric acid the 
compound As.,(S0 4 ) 3 + SO.,. Arsenious oxide is used m the 
manufacture of arsenical compounds. Its poisonous qualities 
are well known. Two to four grains are a fatal dose, and oven 
one grain is dangerous. 

Exp. 181. — Heat very cautiously a small piece of metallic arsenic in 
an open inclined tube, G-10 nun. in diameter. Arsenious oxide will 



274 THE FIFTH GKOUP. 

condense in the cool part of the tube in small brilliant octahedral crys- 
tals, best viewed with a magnifying glass. 

Exp. 182. — Place arsenious oxide in a narrow hard-glass tube, closed 
at one end, and above the oxide a piece of charcoal. Heat the latter to 
glowing, and, by inclining the tube in the flame, slowly sublime the 
oxide, and at the same time keep the charcoal red hot. 

Exp. 183. — a. Heat rapidly arsenious oxide in a narrow tube closed 
at one end. Note fully all observations, b. Place a little arsenious 
oxide on a watch-glass, cover with another glass, and warm with a 
small flame, taking care to sublime the oxide very slowly. The crys- 
tals which form on the upper watch-glass are best seen under a micro- 
scope. 

Potassium Arsenite. — When an excess of arsenious oxide is 
treated with a solution of potassium hydroxide, an acid potas- 
sium arsenite is formed, and, if alcohol is added to the solu- 
tion, after some days crystals of 2KH(As0 2 ) 2 -f- H 2 separate. 
If this salt is boiled with a solution of potassium carbonate, 
carbon dioxide is evolved, and potassium arsenite, KAs0 2 , is 
formed, which may be precipitated from the thick liquid by 
alcohol. A solution of potassium arsenite, known as Fowler's 
solution, is used in medicine. 

Copper Arsenite, Scheele's Green, CuHAs0 3 , is prepared by 
adding an ammoniacal solution of arsenious oxide to a solution 
of copper sulphate. It has a yellowish-green color. Paris 
Green is chiefly composed of a compound of copper arse- 
nite and acetate, having the composition 3Cu(As0 2 ) 2 + 
Cu(C 2 H 3 2 ) 2 . It is extensively used as a green pigment, and 
also as a poison. It is made by boiling a mixture of verdigris 
(a basic copper acetate), arsenious oxide, and water. 

Arsenic Anhydride or Arsenic Pentoxide, As 2 5 , is prepared 
by heating arsenic acid to dull redness. It dissolves slowly in 
water, with formation of orthoarsenic acid. At high tem- 
peratures it decomposes into arsenious oxide and oxygen. 



ARSENIC TRISULPHIDE. 275 

When metallic arsenic burns in air only arsenious oxide is 
formed, and not the pentoxide, differing in this respect from 
phosphorus, which burns to the pentoxide. 

Arsenic Acid is prepared by digesting arsenious oxide with 
nitric acid or aqua regia. The solution is evaporated to dry- 
ness, and the residue gently heated. If it is dissolved in 
water, and the solution evaporated to a thick liquid and 
cooled, crystals having the composition 2AsO(OH) 3 -\- H 2 
separate. These crystals are converted at 100° into anhydrous 
orthoarsenic acid, AsO(OH) 3 . This at 180° changes to diar- 
senic acid, H 4 As 2 7 , and at 200° metarsenic acid, HAs0 3 , is 
formed. These last two acids dissolve in water, with forma- 
tion of the ortho-acid. The arsenates closely resemble the 
corresponding phosphates. Arsenic acid and its salts are em- 
ployed in dyeing, and in the manufacture of dyes. 

Arsenic Disulphide, As 2 S 2 , occurs native as realgar, having 
a red or orange color. An impure disulphide is manufactured 
by heating together arsenious oxide and sulphur. It is used 
in white Indian fire, which is a mixture of one part of disul- 
phide and twelve parts of nitre. 

Arsenic Trisulphide, As 2 S 3 , occurs as a golden-yellow mineral. 
It forms a bright-yellow precipitate when hydrogen sulphide 
is passed into an acid solution of arsenious oxide. It dissolves 
in a solution of ammonium sulphide, with formation of am- 
monium sulpharsenite, (NH 4 ) 3 AsS 3 . 

Arsenic Pentasulphide, As 2 S 6 .— When hydrogen sulphide is passed into 
a solution of arsenic acid no precipitate appears at first, but the arsenic 
acid is reduced to As 2 3) which is then converted into trisulphide : 

2H 3 As0 4 -f 2II 2 S = As 2 3 -f 5II 2 + 2S, 
As 2 3 + 31I 2 S = AsaS 8 + 3H a O. 

If arsenic trisulphide is digested with a. yellow solution o( sodium sul- 
phide, and the solution evaporated, crystals of sodium sulpharsenate, 



276 THE FIFTH GROUP. 

Na 3 AsS4 + 7i"H 2 0, separate. Dilute acid added to a solution of this 
salt causes a precipitation of arsenic pentasulphide: 

2Na 3 AsS 4 + 6HC1 = As 2 S 5 + 6NaCl + 3H 2 S. 



Antimony (Stibium), So. 

Atomic Weight, 120. Density, 6.7. 

Antimony is found in the free state in small quantities. It 
occurs in a number of minerals, the most important being 
stibnite, Sb 2 S 3 , which is the chief source of the antimony of 
commerce. The metal is obtained either by melting the ore 
with wrought-iron, which combines with the sulphur ; or the 
ore is roasted, whereby the sulphur burns to dioxide, and the 
antimony is converted into oxide, which is reduced by fusion 
with coal. The molten metal after a process of purification is 
poured, together with an easily fusible slag, into a thick iron 
mould. The antimony solidifies before the slag, and assumes 
on its upper surface a peculiar fern-like structure, required in 
commerce as evidence of its purity. 

Antimony is a brilliant white metal, hard, and so brittle 
that it is easily pulverized. It crystallizes in rhombohedrons, 
melts at 425°, and vaporizes at a red heat, but too slowly for 
a determination of its gas density. The metal preserves its 
lustre in air unless heated, when it burns with formation of a 
fume of oxide. It dissolves in hot concentrated hydrochloric 
acid, but better in aqua regia. Antimony combines directly 
with sulphur and the halogens. It is extensively used in the 
arts in alloys, to impart hardness, and in making tartar 
emetic, and a few other medicinal preparations. 

Alloys of Antimony. — Type metal is an alloy of lead, tin, 
and antimony, to which copper is sometimes added. Britannia 
metal or pewter, a cheap alloy used as a basis for silver-plated 



ANTIMONY TRICHLORIDE. 277 

ware, consists of tin with about 10 per cent of antimony. 
Anti-friction metal, used for lining journal-boxes, is made of 
antimony, tin, and a little copper. Sometimes lead is added 
to the alloy. 



Antimony Hydride or Stibine, SbH 3 , is formed under the 
same conditions as its analogue arsine, and is found in hydro- 
gen evolved from acid solutions containing antimony. Stibine 
is best prepared by treating a pulverized alloy of 2 parts of 
antimony and 3 of zinc with dilute sulphuric acid. It crys- 
tallizes at —102°, as a snow-white mass, melting at —91° to a 
colorless liquid, which decomposes partially at low tempera- 
tures (—56° to —65°), with separation of antimony. It boils at 
—18°. The gas is odorless, slightly soluble in water, burns 
with a greenish-white name and a fume of antimony oxide. 
If a piece of cold porcelain is held in a hydrogen flame con- 
taining stibine, a black mirror of antimony is deposited. The 
gas is completely decomposed at a red heat, and metallic anti- 
mony may be obtained as a mirror in a tube after the method 
employed with arsine. 

Antimony Trichloride, SbCl 3 . — A concentrated solution of 
this compound has long been known as butter of antimony. 
It is used as a caustic, and for imparting a brownish tint to 
metallic surfaces. Antimony trichloride is formed when 
chlorine is passed over an excess of the metal, and also when 
antimony or antimony sulphide is dissolved in hydrochloric 
acid. On distilling the solution water and acid first pass over, 
and then the antimony trichloride. 

Antimony trichloride forms a transparent crystalline mass, 
melting at 72°, and boiling at 223°. Its vapor is colorless, 
and lias been found to have a density of 117, theory requiring 
113 for SbCl.,. On pouring an acid solution of the trichloride 
into water a white precipitate forms, known as powder of ul- 



278 THE FIFTH GKOUP. 

garotli, which is a mixture of oxychloride and oxide of anti- 
mony. The precipitation is prevented by tartaric acid, or by a 
large excess of hydrochloric acid. If the trichloride is mixed 
with a small quantity of water, and the excess of the trichlo- 
ride removed by washing with ether, Antimonyl Chloride, 
SbOCl, remains. This, on heating, leaves the oxychloride, 
Sb 4 5 Cl 2 . 

Antimony Pentachloride, SbCl., is formed by the direct 
union of its elements, with evolution of light and heat. It is 
prepared by saturating the fused trichloride with chlorine. 
It is a colorless or yellowish fuming liquid, which partially 
decomposes on distillation into trichloride and chlorine. With 
a little water it forms the hydrate SbCl 5 -\- 4H 2 ; with more 
water the oxychloride Sb0 2 Cl, which is converted in antimonic 
acid by hot water. When mixed with a large excess of water 
a clear solution results. 

The compounds SbF 3 , SbF 5 , SbBr 3 , and Sbl 3 are known. 
The pentabromide and pentiodide do not appear to exist. 



Oxides and Hydroxides of Antimony. 

/OH 

Antimonious anhydride, Sb 4 6 . Orthoantimonious acid, Sb(-OH 

\OH* 
Antimony tetroxide, Sb 2 4 . /OH 

Orthoantimonic acid, SbO^ OH. 

\OH 

Antimonic anhydride, Sb s 5 . /SbO<^^ 

Diantimonic acid, 0<^ 

Metantimonic acid, Sb0 2 -OH. 

Antimonious Oxide, Sb 4 6 , is formed, together with tetrox- 
ide, when antimony burns in air, and is obtained pure by 
washing algaroth powder with hot water, and finally with a 



ANTIMONIC ANHYDRIDE. 279 

solution of sodium carbonate, whereby all the chlorine is re- 
moved. The gas density of antimonious oxide, at a tempera- 
ture of about 1560°, has been found in two experiments to be 
283.6 and 289; theory requires 288 for Sb 4 6 . It is both 
basic and acidic in character. It dissolves in sulphuric acid, 
with formation of Sb 2 (S0 4 ) 3 , and in a boiling solution of 
sodium hydroxide, to form sodium antimonite, SbO.ONa. 
Antimonious oxide dissolves readily in a solution of hydrogen 
potassium tartrate (cream of tartar), with formation of anti- 
mony potassium tartrate (tartar emetic), C 4 H 4 SbK0 7 . This 
compound yields, with dilute nitric acid, orthoantimonious 
acid, Sb(OH) 3 . Antimonious oxide is soluble in a solution of 
tartaric acid. 

Antimony Tetroxide, Sb 2 4 . — This oxide is formed when 
antimonic acid is strongly heated, and is prepared by oxidiz- 
ing antimony with concentrated nitric acid and igniting the 
residue. It is infusible and non-volatile, and is dissolved by 
hydrochloric acid with great difficulty. 

Antimonic Anhydride or Antimony Pentoxide, S 2 5 , is ob- 
tained by treating antimony with an excess of strong nitric 
acid, evaporating off the acid, and then heating the residue 
to a temperature below redness. It is a pale-yellow powder, 
practically insoluble in water and nitric acid, but slowly sol- 
uble in hydrochloric acid. It reddens moist litmus paper. 

Antimonic Acid. — When a solution of a nietantimonate is treated with 
sulphuric or nitric acid a white powder separates, which, according to 
some authorities, has the composition ShO(OH) 3 , and according to 
others Sb(OH) 5 . It loses water at 175°, and is converted into metanti- 
monic acid, Sb0 2 .OH. The potassium salt of this acid is prepared by 
fusing a mixture of 1 pail of antimony and 4 parts of potassium nitrate. 
The mass is treated with tepid water to remove potassium nitrite, and 
the white powder obtained is boiled for an hour or two with water to 
dissolve it. The solution on evaporation leaves a gummy mass, having 
the composition 2Ki5b0 3 + 51I.JO, and soluble in warm water. 



280 THE FIFTH GROUP. 

Potassium Diantimonate, K4SD-2O7, is prepared by fusing potassium 
antimonate with three times its weight of potassium hydroxide, dis- 
solving the product in water, and evaporating the solution. Deliques- 
cent warty crystals separate, which are decomposed by water into free 
alkali and hydrogen potassium diantimonate, H 2 K2.Sb 2 07 + 6H 2 0. 
This last salt is easily soluble in water at 40° to 5CT, and in solution de- 
composes into gelatinous antimonate. 

Hydrogen Sodium Diantimonate, H 2 Na 2 Sb 2 7 -f 6H 2 0. — This 
sodium salt is remarkable for its insolubility, sodium salts as a 
rule being soluble. It separates as a crystalline precipitate 
when hydrogen potassium diantimonate is added to a solution 
of a sodium salt. 

It will be observed that the metantimonates stand in the 
same relation to the diantimonates as the metaphosphates to 
the diphosphates, and hence are thus named. The metanti- 
monates are, however, often termed antimonates, and the di- 
antimonates metantimonates. 

Antimony Trisulphide, Sb 2 S 3 , occurs native in lead-gray 
crystals of stibnite. Hydrogen sulphide produces in an acid 
solution of an antimonous salt an orange-red amorphous pre- 
cipitate of antimony trisulphide containing water. The pre- 
cipitate loses water on heating, and is converted at 200° into 
the black crystalline form. Antimony trisulphide fuses easily, 
and may be volatilized in an atmosphere of nitrogen, It dis- 
solves in hydrochloric acid, with evolution of hydrogen sul- 
phide. It is soluble in solutions of potassium or sodium 
hydroxide and alkali sulphides. Antimony sulphide is em- 
ployed in the preparation of other antimony compounds, and 
is used for vulcanizing rubber. 

Antimony Pentasulphide, Sb 2 S 5 , is prepared by passing hydrogen sul- 
phide through a solution of the pentachloride containing tartaric acid, 
and also by decomposing a solution of a sulphantimonate with an acid. 
It is a yellowish-red powder, soluble in alkalies and alkali sulphides. It 
is decomposed by heat into trisulphide and sulphur. It is a sulphur- 



BISMUTH. 281 

acid anhydride, uniting with metallic sulphides to form sulphosalts, 
most of which have the general formula R 3 SbS 4 , analogous to the ortho- 
phosphates. 

Sodium Sulphantimonate, Na 3 SbS 4 -j-9H 2 0, is formed when either of the 
sulphides of antimony is dissolved in a solution of sodium sulphide, and 
also when a mixture of antimouious oxide and sulphur is boiled with a 
solution of sodium hydroxide. The solution is evaporated until crystals 
form on cooling, and is then allowed to cool slowly. The crystals are 
colorless or pale yellow, and are soluble in 2.9 parts of water at 15°. 



Bismuth; Bi. 

Atomic Weight, 208. Density, 9.8. 

The chief source of this comparatively rare metal is native 
bismuth. Bismuth has a reddish-white color, a brilliant 
lustre, and is highly crystalline. It melts at 270°, and vola- 
tilizes at very high temperatures. It dissolves readily in nitric 
acid and aqua regia. 

Its salts are used in medicine, and the metal itself enters 
into the composition of various fusible alloys. Wood's fusible 
metal, melting at 60°. 5, is composed of 4 parts of bismuth, 
2 of lead, 1 of tin, and. 1 of cadmium. Fusible alloys, melt- 
ing at certain temperatures, are used for safety plugs in steam- 
boilers, and when, from increase of pressure, the temperature 
of the steam reaches the melting point of the alloy, the plug 
opens and allows steam to escape. 

Fusible metal is valuable for making casts, on account of 
its low melting point, and its property of expanding in solidi- 
fying, thus making a good cast. 

Bismuth in chemical properties is more basic and more 
metallic in character than antimony. Its halides resemble 
the halides of antimony and phosphorus in that they are de- 
composed by water into oxy-compounds. No bismuth com- 
pound of hydrogen is known. 



282 THE FIFTH GEOUP. 

Bismuth Dichloride, BiCl 2 , is obtained by beating the tri- 
chloride with metallic bismuth, and by other processes. Its 
molecular weight is not known, as it decomposes on heating 
into the trichloride and bismuth. 

Bismuth Trichloride, BiCl 3 , is formed when the metal is 
burned in chlorine. It is a white fusible substance. Its gas 
density has been found to accord with the formula BiCl 3 . 
Water decomposes it, with formation of insoluble oxychloride 
or bismuthyl chloride, BiOCl, thus : 

BiCl 3 + H 2 = BiOCl -f 2HC1. 

No pentachloride of bismuth is known. There are bromine, 
iodine, and fluorine analogues of the trichloride. 

Bismuth Dioxide, Bi 2 2 , is a grayish-black powder, which 
burns when heated in air to the trioxide. 

Bismuth Trioxide, Bi 2 3 , occurs as bismuth ochre. It is 
formed when the metal is burned in air, and by igniting the 
hydroxide or nitrate. 

Bismuth Trihydroxide, Bi(0H) 3 , separates when a solution 
of bismuth nitrate is dropped into a cold solution of potassium 
hydroxide. On drying at 100° it is converted into bismuthyl 
hydroxide, BiO.OH. 

N0„-(k 

Bismuth Nitrate, Bi(N0 3 ) 3 or N0 2 -0~Bi.— This salt is pre- 

N0 2 -(K 

pared by dissolving the metal in nitric acid. The solution 
yields on evaporation hydrous crystals. 

N0 2 -(k 

Basic Bismuth Nitrate, HO-^Bi.— This basic salt is ob- 



VANADIUM. 283 

tained as a curdy mass when bismuth nitrate is treated with 
cold water. If this remains in contact with the acid solution 
it is gradually converted into the salt 5Bi 2 3 .41Sr 2 5 + 9H 2 0. 
On pouring a nitric acid solution of bismuth into water a 
basic salt separates., varying somewhat in composition accord- 
ing to the amount of water used in precipitating and washing. 
It has been long used in medicine, and is the most common 
preparation of bismuth. 

Bismuth Sulphate, Bi 2 (S0 4 )3, is said to remain when a solution of bis- 
muth or its oxide in concentrated sulphuric acid is evaporated. The 
residue yields on treatment with water a basic salt, which loses water 
on heating, with formation of bismuthyl sulphate, (BiO) 2 S0 4 . 

Bismuth Tetroxide, Bi 2 4 , Bismuth Pentoxide, Bi 2 5 , and Bismuthic 
Acid, HBi0 3 , have been prepared. 

Bismuth Trisulphide, Bi 2 S 3 , occurs as bismuthite. It can be prepared 
by melting together sulphur and bismuth, and also by precipitating a 
solution of bismuth with hydrogen sulphide. Bismuth Disulphide, Bi 2 S 2 , 
is known. 



Yanadium, Y. 

Atomic Weight, 51.3. Density, 5.5. 

In 1801 Del Rio found a new element in a Mexican lead ore. Later, 
it was regarded as an impure chromium oxide. In 1830 Sefstrom dis- 
covered the same element in the ores of Taberg, and named it Vana- 
dium, after Vanades, a Scandinavian goddess. Vanadium occurs very 
sparingly, though widely distributed, having been found in a number 
of clays and iron ores, and in trap and basalt, The metal is obtained 
with difficulty by reducing the dichloride at a red heat with hydrogen 
absolutely free from traces of oxygen Metallic vanadium thus pre- 
pared is a light whitish-gray powder, which appears as a brilliant white 
crystalline mass under the microscope. It oxidizes slowly in the air, 
and does not decompose water at common temperatures. It burns bril- 
liantly in the flame or when heated in oxygen. It does not fuse at a 
red heat in hydrogen. It is insoluble in hydrochloric acid ami cold 



284 THE FIFTH GROUP. 

sulphuric acid, whilst it dissolves quickly in nitric acid, forming a blue 
solution. 

Vanadium Chlorides. — Vanadium Tetrachloride, VC1 4 , is a brownish-red 
liquid, boiling at 154°, with partial decomposition. Its observed gas 
density is 96.7. Vanadium Trichloride, VC1 3 , and Vanadium Dichloride, 
VClo, are formed when the foregoing compound is passed with hydrogen 
through a tube heated to dull redness. At a higher temperature they 
are reduced to the metal. No vanadium penta chloride is known. 

Vanadium Oxychlorides.— Vanadyl Trichloride, V0C1 3 , corresponds to 
phosphorus oxychioride. It is a lemon-yellow liquid, boiling at 126°. 7, 
and having a gas density of 88.2. Vanadyl Dichloride, V0C1 2 , Vanadyl 
Monochloride, V0C1, and Divanadyl Monochloride, V 2 2 C1, are formed 
when vanadyl trichloride is reduced by hydrogen at a red heat. 

No fluoride or oxynuoride of vanadium is known, but a number of 
fluoxyvanadates have been prepared, one of which is 6NH 4 F.V 2 5 . 
2VOF 3 + 2H 2 0. 

Vanadium Oxides are analogous to the oxides of nitrogen : 

Vanadium monoxide, V 2 

Vanadium dioxide or hypovanadious oxide, V 2 2 
Vanadium trioxide or vanadious oxide, . . V 2 3 
Vanadium tetroxide or hypovanadic oxide, V 2 4 
Vanadium pentoxide or vanadic oxide, . . V 2 5 

No salts of vanadium monoxide have been prepared. All the others 
form salts with acids, and the two highest oxides are acid-forming 
oxides. The monoxide is formed by the slow oxidation of- the metal, 
while the trioxide and tetroxide are produced by the partial reduction 
of the pentoxide. 

Vanadium Dioxide, V 2 2 , may be obtained by reducing the higher ox- 
ides by potassium. It was regarded as metallic vanadium by Berzelius. 
It conducts electricity, and burns when heated in air to the trioxide. 
It is insoluble in acids, except aqua regia, 

Vanadium Pentoxide, V 2 5 , is the most important of the vanadium 
oxides. It is obtained by roasting ammonium meta vanadate, and in a 
purer state by decomposing vanadyl trichloride by water. Vanadium 
pentoxide is soluble in about 1000 parts of water, forming a tasteless 
yellow solution which turns blue litmus paper red. It is both a basic 
and an acidic oxide. With strong acids it forms salts of the trivalent 
radical vanadyl, VO. 

Vanadyl Sulphate, (V0) 2 (S0 4 ) 3 , is obtained by dissolving the pentoxide 
in hot sulphuric acid. 



NIOBIUM. 285 

Orthovanadic Acid, VO(OH) 3 , is not known. 

Divanadic Acid, H 4 V 2 7 , has been prepared. 

Metavanadic Acid, V0 2 OH, is a yellow pigment known as vanadium 
bronze. It is obtained as brilliant golden scales by boiling copper 
vanadate with a solution of sulphurous acid. 

The vanadates correspond to the phosphates. Sodium Metavanadate, 
NaV0 3 ; Sodium Or tho vanadate, Na 3 V0 4 ; Sodium Divanadate, Na 4 V 2 7 ; 
and more complex poly vanadates, have been described. 

Ammonium Metavanadate, NH 4 V0 3 , is prepared by dissolving vanadium 
pentoxide in ammonia and evaporating .the solution. It is insoluble in 
solution of ammonium chloride, and separates when solid ammonium 
chloride is left in a solution of sodium meta- or di- vanadate : 

Na 4 V 9 7 +4NH 4 Cl = 2NH 4 V0 3 + 2NH 3 + H 2 + 4NaCl. 

Vanadium Nitrides. — Vanadium is one of the few elements which 
unite directly with nitrogen. The mononitride, VN, is formed when 
the metal is heated in nitrogen, and also when the oxide is heated in 
ammonia. Berzelius mistook the vanadium nitride for the metal. The 
dinitride, VN 2 , is known. 



Niobium, Nb. 

Atomic Weight, 94. Density, 7. 

Niobium is found in columbite and a few other rare minerals, and is 
usually associated with tantalum. The metal is obtained in the form 
of a steel-gray crust when the vapor of the chloride mixed with hydro- 
gen is passed through a strongly heated tube. It ignites at a low tem- 
perature and burns brilliantly in air. The metal is soluble in sulphuric 
acid, but is only slightly attacked by aqua regia. 

Niobium Trichloride, NbCl 3 , is obtained bypassing the vapor of the 
pentachloride through a heated tube. It is non-volatile, and is not de- 
composed by water. When heated in carbon dioxide, the oxychloride, 
NbOCla, results. This reaction is not exhibited by any other metallic 
chloride. (Roscoe.) 

Niobium Pentachloride, NbCl r ,, forms in yellow needles when chlorine 
gas is passed through a heated mixture of the pentoxide and charcoal. 
It fuses at 194°, and boils at 240°. 5. Its observed gas density is 188°.6, 
the formula NbCl 6 requiring lo5.3. 



280 THE FIFTH GROUP. 

Niobium Oxychloride or Niobyl Chloride, NbOCl 3 , is formed by the direct 
union of niobium dioxide with chlorine. Its observed gas density is 
113, theory requiring 108. 

Niobium Oxyfluoride, NbOF 3 , has been prepared. The pentafluoride, 
NbF 5 , is known only in combination, e.g., potassium niobium fluoride, 
2KF.XbF 5 . 

Niobium Dioxide, Nb 2 2 , is obtained by intensely heating a mixture of 
potassium oxyfluoride, 3KF.NbOF 3 , and sodium. It is an insoluble 
white powder, which burns brilliantly when heated in air. 

Niobium Tetroxide, Nb 2 4 , is obtained in the form of a heavy black 
powder when the pentoxide is strongly ignited in hydrogeu. It is in- 
soluble in acids. 

Niobium Pentoxide, Nb 2 5 , is prepared by heating potassium niobium 
fluoride with sulphuric acid, and dissolving out the potassium sulphate 
formed with water. The residue is then heated with ammonium car- 
bonate to remove sulphuric acid. Thus obtained it is a white amor- 
phous powder, which becomes denser and crystalline on ignition. 

Niobium Hydroxide or Niobic Acid is formed when the pentachloride 
or oxychloride is decomposed by water. The product, dried at 100 c , 
contains varying proportions of water. Niobic acid is soluble in alkali 
carbonates and hydroxides, and after treatment with hot hydrochloric 
acid is soluble in water. 

Mobic acid forms complex poly-acid salts, of which the best known 
is potassium hexniobate, KJSTbeOig + 16H 2 0. This salt is formed when 
niobium pentoxide is fused with twice its weight of potassium carbonate. 



Didyniium, Di. 

Atomic Weight, 145. Density, 6.54. 

This element occurs very sparingly, aud is associated with cerium 
and lanthanum. Metallic didymium is obtained by the electrolysis of 
the chloride. It is ductile, has an almost white color, and tarnishes in 
air, and in the form of filings burns brilliantly in the flame. The 
didymium compounds closely resemble the lanthanum compounds, and 
in general present more analogy to the compounds of the third group 
than to those of the fifth. Its highest oxide. Di 2 5 , is the analogue of 
the highest oxides of the members of the fifth group. 



SAMARIUM. 28? 

Didymium Chloride, DiCl 3 + 6H 2 0, crystallizes from a solution of the 
oxide in hydrochloric acid. 

Didymium Sesquioxide, Di 2 3 , forms a white powder when the nitrate 
or oxalate is ignited. It is a strongly basic oxide, dissolving in dilute 
acids, and setting ammonia free from ammonium salts. It absorbs 
carbon dioxide from the air. In hot water it changes gradually to the 
hydroxide. 

Potassium hydroxide precipitates the hydroxide Di(OH) 3 from a solu- 
tion of the chloride. 

Didymium Sulphate, Di 2 (S0 4 ) 3 + 8H 2 0, is converted by ignition at a red 
heat into (DiO) 2 S0 4 , the analogue of vanadyl and bismuthyl sulphates. 

Didymium Pentoxide, Di 2 5 , results from the ignition of the nitrate 
(Bruner). It dissolves in dilute sulphuric or nitric acid, without evolu- 
tion of gas, but in concentrated acids ozonized oxygen is set free. On 
adding a solution of potassium hydroxide to a solution of didymium 
nitrate and hydrogen dioxide, the hydroxide DiO(OH) 3 separates. This 
compound, analogous in formula to orthophosphoric acid, does not 
form salts with bases. 



Samarium, Sm. 

Atomic Weight, 150. 

This very rare element was discovered by means of the spectroscope 
in 1879, by Lecoc de Boisbaudran, in the mineral samarskite. It has 
been found in a few other minerals. The metal has not been isolated. 
Samarium compounds resemble those of didymium, and, like the latter, 
samarium exhibits analogy to the elements of the third group. 

Samarium Chloride, SmCl 3 -j- 6H 2 0, forms large deliquescent crystals. 

Samarium Hydroxide is white, gelatinous, and insoluble in alkalies. 
It is converted by ignition into the oxide Sm 2 O s . A higher oxide is 
precipitated when a solution of samarium and hydrogen dioxide is 
treated with ammonia. The precipitate loses oxygen readily, and after 
drying has the composition Sm 4 9 , or very nearly Sm 2 6 . 

Samarium Sulphate, Sm 2 (S0 4 ) 3 -j- 8H 2 0, forms brilliant crystals. It does 
not form an alum, with alkali sulphates. 



288 THE FIFTH GROUP. 

Tantalum, Ta. 

Atomic Weight, 183. 

Metallic tantalum apparently has not been prepared in the pure state. 
Berzelius, by heating potassium tantalum fluoride with potassium, ob- 
tained a black powder which burned to oxide when heated in air, and 
was insoluble in acids, excepting hydrofluoric. The tantalum com- 
pounds resemble those of niobium. 

Tantalum Pentachloride, TaCl s , is obtained by the same method as 
niobium pentachloride. It melts at 211°, and boils at 240°. Its observed 
gas density is 183.8, the formula TaCl 5 requiring 180. 

Tantalum Pentafluoride, TaF 5 , is known only in solution, or in com- 
bination with other fluorides, as for example, 2KF.TaF 5 . 

Tantalum Tetroxide, Ta 2 4 , results from intense ignition of the pent- 
oxide in a carbon crucible. It is very hard, insoluble in all acids, and 
when heated burns to the pentoxide. 

Tantalum Pentoxide, Ta 2 5 , is obtained by the method described for 
the preparation of niobium pentoxide. It is insoluble in acids, and has 
a density of 7.35-8 

Tantalum Hydroxide or Tantalic Acid, HTa0 3 , is obtained as a crys- 
talline powder when tantalum pentachloride is decomposed by moist 
air, and the product treated with ammonia, and dried at 100°. It is 
soluble in hydrogen potassium oxalate, and hydrofluoric acid. It is 
converted into the pentoxide by ignition. 

Tantalates. — Some native tantalates are normal salts of the above 
acid. Sodium and potassium salts of the unknown acid, H 8 Ta 6 0i9, 
and also other polytantalates, have been prepared. 



Summary of the Fifth Group. 

Hydrides axd Halides. 



Nitrogen, . . 


. N 2 H 4 


NH 3 




NC1 3 


Phosphorus, . 


. P 2 H 4 


PH 3 


PI 2 


PCI 3 


Arsenic, . . 




AsH 3 


Asl 2 


AsCl 3 


Antimony, 




SbH 3 




SbCl 3 


Bismuth, . . 






BiCl 2 


BiCl 3 


Vanadium, 






VC1 2 


VC1 3 


Niobium, . . 








NbCl 3 



pels 

Asl 3 

SbCl 5 

VCL 

NbCL 






SUMMARY OF THE FIFTH GEOUP. 



289 



Didymium DiCl 3 

Samarium SmCl 3 

Tantalum, TaCl 6 

Oxides. 

Nitrogen, N 2 NO N 2 3 N 2 4 N 2 5 

Phosphorus, P 2 3 P 2 4 P 2 5 

Arsenic, As 4 6 As 2 5 

Antimony, . . . . ; Sb 4 6 Sb 2 4 Sb 2 5 

Bismuth, Bi 2 2 Bi 2 3 Bi 2 4 Bi 2 5 

Vanadium, V 2 V 2 2 V 2 3 V 2 4 V 2 5 

Niobium, ....... Nb 2 2 Nb 2 4 Nb 2 5 

Didymium Di 2 3 Di 2 5 

Samarium Sm 2 3 Sm 2 5 ? 

Tantalum, Ta 2 4 Ta 2 5 

Hydroxides. 

Nitrogen, .NOH? NO.OH N0 2 .OH 

Phosphorus, . P(OH) 3 PO(OH) 3 (PO) 2 0(OH) 4 P0 2 .OH 

Arsenic, . . As(OH) 3 ? AsO(OH) 3 (AsO) 2 0(OH) 4 As0 2 .OH 

Antimony, . Sb(OH) 3 SbO(OH) 3 (SbO) 2 0(OH) 4 Sb0 2 .OH 

Bismuth, . . Bi(OH) 3 Bi0 2 .OH 

Vanadium, . (VO) 2 0(OH) 4 V0 2 .OH 

Niobium, . " Composition of hydroxide unknown. 

Didymium . Di(OH) 3 DiO(OH) 3 

Samarium . Composition of hydroxides unknown. 

Tantalum, . Ta0 2 .OH 

The distinguishing characteristics of the fifth group are 
exhibited in the pentoxides, which are all acid anhydrides, 
excepting those of didymium and samarium, and in the 
acid hydroxides, of which nitric acid, NO„.OH, is the first 
member. No other group of elements forms similar classes 
of compounds. The trichlorides and trioxides indicate a 
close relationship in chemical properties, but are not charac- 
teristic of the fifth group, showing rather an analogy in for- 
mula to the trichlorides and oxides of the third group, of 
which AlOlg and A1 9 3 are examples. 

The hydrides, NH 3 , PH S , AsII 3 , and SbH 3 , exhibit a gradation 
of properties with increasing atomic weights. Nitrogen, with 



290 



THE FIFTH GROUP. 



the lowest atomic weight, forms with hydrogen a stable basic 
compound ; phosphine, PH 3 , is less stable and but feebly 
basic ; while arsine, AsH 3 , and stibine, SbH 3 , are still less 
stable, and do not combine with acids. 

Nitrogen stands apart from all the other members of the 
group in many properties. This is seen in ammonia, and 
in the nitrogen halides, which are highly explosive, and do 
not react with water to form oxides or hydroxides as do the 
tri- and pentachlorides of the other elements of the group, 
VaClg excepted. Nitrogen and phosphorus do not form basic 
oxides and hydroxides, and there are no salts in which these 
elements are basic radicals. Arsenic, antimony, and bismuth 
form salts with oxygen acids, and thus exhibit basic char- 
acters. Their sulphates are As 2 (S0 4 ) 3 + S0 3 , Sb 2 (S0 4 ) 3 , 
Bi 2 (S0 4 ) 3 . They are decomposed by water into basic salts. 
Vanadium, on the other hand, forms vanadyl sulphate, 
(VO) 2 (S0 4 ) 3 , while well-defined sulphates of niobium and 
tantalum are unknown. 

The densities of the elements of the fifth group increase 
with their atomic weights. Nitrogen and phosphorus, like 
other non-metals, have low densities. Their melting and 
boiling points are widely separated from the' melting and 
boiling points of the remaining members of the group. 
Nitrogen and, to a less degree, phosphorus stand apart in 
physical as well as in chemical properties. 





Atomic 




Gas 


Melting 


Boiling 




W x eight. 


Density. 


Density 


Point. 


Point. 


Nitrogen, . . 


14 


0.88 (liquid) 


14 


-214° 


-194° 


Phosphorus, . 


31 


1.83 


62 


44° 


278° 


Vanadium, . 


51.3 


5.5 


— 


above red heat. 





Arsenic, . . . 


75 - 


5.7 


150 


dull red. 





Niobium, . . 


94 


7 


— 


abo^e red heat. 





Antimony, . . 


120 


6.7 


— 


425° 





Didymium, . 


145 


6.5 


— 








Samarium, 


. 150 


— 


— 








Tantalum, 


183 


— 


— 





— - 


Bismuth, . . 


208 . 


9.8 


— 


270° 


- — - 



THE THIED GROUP. 

The members of this group are boron and aluminum, and 
the rare elements gallium, indium, thallium, scandium, 
yttrium, lanthanum, erbium, and ytterbium. 

They are all trivalent elements, although thallium is uni- 
valent in thallous compounds, and boron pentavalent in a few 
compounds. 



Boron, B. 

Atomic Weight, 11. Density, 2.7. 

This element is found in native borates and a few other 
minerals. In order to obtain free boron a mixture of boric 
oxide and sodium is intensely heated under a layer of common 
salt, and the molten mass is poured into dilute hydrochloric 
acid. The sodium chloride, sodium borate, and boric oxide 
dissolve, and the boron remains as an amorphous brown pow- 
der. Amorphous boron does not oxidize at common tempera- 
tures, but burns brilliantly to boric oxide when heated in air. 
It dissolves in molten aluminum, and when the mass cools 
brownish -yellow crystals form. These may be separated by 
dissolving the aluminum in hydrochloric acid or a solution of 
sodium hydroxide, and further purified by treatment with a 
mixture of nitric and hydrofluoric acids. The crystals are not 
pure boron, but contain a little aluminum, the proportion of 
the two elements varying in different preparations. The 
analysis of one sample correspon (led to the formula A1B 1S , and 
of another to A1B., 3 . Pure crystalline boron dors not appear to 



292 THE THIRD GROUP. 

have been obtained. The above-described crystals are, how- 
ever, regarded as a modification of boron. They have a 
brilliant lustre, are almost as hard as the diamond, and will 
scratch corundum. They are insoluble in acids, but are acted 
on by molten sodium hydroxide, with formation of sodium 
borate. In the preparation of crystalline boron, graphite-like 
scales, of the composition iUB 2 , have occasionally been ob- 
tained. These were formerly regarded as a third modification 
of boron. 



Boron Hydride, BH3 1 ? — When magnesium dust is fused with boric oxide 
an impure magnesium boride is formed. This yields on treating with 
hydrochloric acid a gas which is mostly h} r drogen, but which contains 
a volatile boron compound, probably having the composition BH 3 . 

/ F 
Boron Fluoride, BF 3 or B(— F. — This compound is easily obtained by the 

\F 
following reaction : 

3Ca<! + >0 + 3S0 3 <3ii = 3S0 2 <^>Ca+3H 2 + 2B^F. 

It is a colorless suffocating gas, having a density corresponding to its 
formula. It unites with ammonia to form a white compound having 
the composition BF 3 .lSrH3. This body is dissociated by heat into a 
mixture of the gases BF 3 and NH 3) which combine again on cooling. 



/CI 

Boron Chloride, BC1 3 or B^ CI, is formed by the direct 

\C1 
union of its elements, and also by heating a mixture of boric 
oxide and charcoal in chlorine. It is a colorless liquid, boiling 
at 17°. Its gas density corresponds to its formula. It fumes 
in moist air, and is decomposed by water into hydrochloric and 
boric acids. BBr 3 is known, but the iodide has not been 
obtained. 






BOEOK. 293 



d 



Boron Oxychloride, BOCL or = B^-C1. — This substance 

M3I 

has been obtained as a yellowish green liquid in the preparation 
of boron chloride. It is unchanged at 100°, but decomposes 
at higher temperatures as follows : 

3B0C1 3 = B 2 3 + BCI3 + 3C1 2 . 

It is decomposed slowly by water into boric and Hydrochloric 
acids and free chlorine. This oxychloride is of theoretical 
interest, as it indicates a pentavalent character in boron. 

B=0 

Boric Oxide, Boric Anhydride, B 2 3 or >0. — This is the 

B = 

only known oxide of boron. It is easily obtained by heating 
boric acid until the mass is in a quiet state of fusion. On 
cooling it cracks spontaneously. It is soluble in water, and 
becomes opaque in moist air. At high temperatures boric 
oxide decomposes salts of acids more volatile than itself. 

/OH 

Boric or Orthoboric Acid, H 3 B0, or B^-OH.— This acid is 

M)H 

formed from boric oxide and water, thus : 

B=0 /OH 

>0 + 3H 2 = 2Bf OH. 

B=0 \OH 

It occurs in volcanic regions in the jets of steam which 
issue from the vents about the volcanoes. Tuscany has sup- 
plied the European market with boric acid. It is also ob- 
tained from native borates. It is soluble in three parts of 
water at 100°, and in about 25 parts at ordinary temperatures. 
When its solution is boiled, a little of the acid passes oil with 
the steam — a fact which explains the presence of boric acid in 



294 THE THIKD GROUP. 

Yolcanic vapors. Alcoholic vapors also carry it, and burn with 
a characteristic green flame. If a piece of turmeric paper is 
moistened with a solution of boric acid and then dried, a pecu- 
liar red coloration appears, which differs from the alkaline 
reaction of turmeric paper in that it is not changed to yellow 
by dilute hydrochloric acid. 

Exp. 184. — a. Place in a test-tube 10 cc. of cold water and 1 cc. of 
concentrated sulphuric acid, beat to boiling, and add 5 grams of borax. 
After tbe latter bas dissolved, allow tbe solution to cool. Pour oft' tbe 
liquid as completely as possible, and wasb tbe crystals twice witb a 
little water. Tbe boric acid obtained will not be free from sulpburic 
acid and sodium sulpbate, but will answer for further experiments. 
Represent by an equation the reaction by which the boric acid is 
formed. 

b. Dissolve a little boric acid in water in a test-tube, boil the solution, 
and hold a Bunsen flame over the escaping vapor. The greenish tinge 
of the gas flame is due to boric acid in the steam. 

c. Dissolve some boric acid in a small porcelain dish in a tablespoon- 
ful of ethyl alcohol, or better, impure methyl alcohol of commerce. 
Warm gently and ignite the alcohol. 

Metaboric Acid, HB0 2 or = B0-H, is formed from ortho- 
boric acid by the removal of one molecule of water : 

™>B-OH = 0=B-OH + H 2 0. 

It remains as a white powder when orthoboric acid is heated 
to 100°. At this temperature, according to Schaffgotsch, it 
volatilizes slowly but completely. 

Tetraboric Acid, B 4 0.(0H) 2 , is a brittle mass attained by 
prolonged heating of metaboric acid to 160°. It may be 
viewed as a derivative of the ortho-acid, thus : 

4H ! B0 3 = B 4 5 (OH), + 5H 3 ; or 





BORON. 


/OH 
B^-OH 


/OH 


•b/ 


\OH 


/ \ 


OH 


u \ / 


b^oh 


\b/ 


\OH 


\ 




\0 + 5H a 0. 


/OH 
Bf-OH 


\OH 


/ \ 


B^-OH 


W 


\0H 


\0H 



295 



Borax, Sodium Tetraborate, Na 2 B 4 7 -f 10 2 H 2 0. — Borax is 
the most important salt of boric acid. It occurs in crystals 
in California, Peru, and in various localities in Asia. Borax is 
found in abundance in some salt lakes. A considerable por- 
tion of the borax of commerce is manufactured from boric 
acid by treatment with sodium carbonate. The pentahydrated 
salt Na 2 B 4 7 -f 5H 2 separates in octahedral crystals, when a 
hot saturated solution cools ; after the temperature falls to 
56° large monoclinic prisms of the decahydrated salt are de- 
posited. Both hydrates are sold under the name of borax. 
The decahydrated salt dissolves in 14 parts of water at or- 
dinary temperature, and in half its weight in water at 100°. 
Borax has a slightly alkaline taste and reaction. It appears 
to separate into boric acid and sodium hydroxide when dis- 
solved in a large quantity of water. If, for example, to a con- 
centrated solution of borax containing litmus acetic acid is 
added until a faint red tint appears, and then much water is 
added, the blue color of litmus is seen. The cleansing action 
of a solution of borax is due in part to the free alkali it 
contains. 

Both the octahedral and prismatic borax swell when 
heated, give off water, and are converted into a white 



296 THE THIRD GROUP. 

porous mass of anhydrous borax, Na 2 B 4 7 , which melts at a 
high temperature, forming borax glass. Molten borax possesses 
the property of dissolving metallic oxides, many of which give 
characteristic colors to the flux. Hence the use of borax in 
blowpipe tests. Borax is used in soldering gold, silver, and 
copper, and in welding iron, to cleanse the surfaces of the 
metals. 

Exp 185. — Add to a concentrated solution of borax a little litmus 
solution and acetic acid to faint acid reaction, then dilute the solution 
largely with water. 

Exp 186. — Heat a small ring on the end of a copper wire above the 
lamp flame until a coating of oxide forms, and fuse on it a little borax, 
taking care to heat in the upper and oxidizing part of the flame. The 
copper oxide will dissolve and color the borax glass green, leaving the 
surface of the copper bright. 

Sodium Metaborate, BO-0-Na, is obtained by fusing to- 
gether equivalent weights of borax and of sodium carbon- 
ate. It crystallizes with 4 molecules of water. The salt 
dry and in solution, absorbs carbon dioxide with formation 
of sodium carbonate and sodium tetraborate. 

/O-Na 

Sodium Orthoborate, Br- O-Na — This salt is formed when 
\0-Ka 

boric oxide is fused with an excess of sodium hydroxide. It 
is converted by water into the metaborate. 






Aluminum, Al. 

Atomic Weight, 27. Density, 2.6. 

Aluminum is the most abundant of the elements excepting 
oxygen and silicon. It is a constituent of soils, most rocks, and 



ALUMINUM. 297 

many minerals. The metal is obtained from its chloride by 
electrolysis or by fusion with sodium. It is manufactured by 
melting together a double chloride of aluminum and sodium, 
metallic sodium and cryolite. The last is added as a flux. 

Aluminum is not reduced from its oxide by carbon or 
hydrogen at the highest temperature of a wind furnace. 
Aluminum oxide, mixed with charcoal, is however reduced at 
the temperature of the voltaic arc, the metal at the same time 
volatilizing. When the reduction is made in the electrical 
furnace part of the vapor condenses in the layer of charcoal 
and part escapes from the furnace with the gases. If, how- 
ever, the reduction is made in the presence of iron or copper 
these metals retain the aluminum, forming alloys of great in- 
dustrial importance. 

Pure aluminum is almost as white as tin, is capable of a 
good polish, and does not oxidize in air. The impure metal of 
commerce, containing small quantities of iron and silicon, 
soon tarnishes owing to formation of a thin coating of oxide. 
Aluminum is ductile, and very malleable. It melts at about 
700°. The cast metal is about as hard as silver, but the ham- 
mered metal has the hardness of soft iron. Aluminum does 
not burn in air or oxygen except in the form of filings, thin 
foil, or fine wire. It dissolves in hydrochloric and sulphuric 
acids, and in solutions of the fixed alkalies. Nitric acid, con- 
centrated or dilute, does not act on it. Organic acids have 
little action, but if common salt is present the metal slowly 
dissolves — a fact which precludes its use in culinary articles. 

It is said that aluminum may be soldered by an alloy con- 
sisting of 5 parts of tin and 1 of aluminum. The metal on 
account of its lightness and great tensile strength is valuable 
for many purposes, but the high cost prevents the extensive 
use of it. 

Aluminum bronze is an alloy .of copper containing about 10 
per cent of aluminum. It has the color of gold, is malleable, 
makes good castings, and is remarkable for its strength and 



298 THE THIRD GROUP. 

permanence in air. An increase of 2 or 3 per cent of 
aluminum renders the alloy brittle, and with increasing pro- 
portions of aluminum the color of the alloy becomes lighter. 

Iron takes up several per cent of aluminum, but not as 
much as copper. The alloy of iron and aluminum is added 
to molten iron, the presence of a fraction of a per cent of 
aluminum having been found to greatly improve the quality 
of castings for various purposes. 

Valence of Aluminum. — The gas density of aluminum 
chloride has been found at 350° to 440° to be 133.4 ; theory 
requires for A1 2 C1 C 128.2. The gas densities of the bromide 
and iodide correspond to the molecules Al 2 Br 6 and A1 2 I 6 . These 
results have led to the view that the double atom Al 2 is tetra- 

. Al— 
valent, having the constitution i ^, and that aluminum is 

tetravalent. Kecently (1887) Nilson has, however, found that 
the density of aluminum chloride diminishes at temperatures 
above 440°, and at 835° to 943° is 65.7 ; calculated for A1C1 3 , 
66.7. Aluminum is therefore tri valent. 



/ C1 

Aluminum Chloride, A1CL or A1(-C1. — A solution of alu- 

\C1 
minum chloride is obtained by dissolving aluminum hydrox- 
ide in hydrochloric acid. If the solution is evaporated at ordi- 
nary temperature, crystals of A1C1 3 + 6H 2 separate. These 
on heating lose hydrochloric acid and water, with the formation 
of alumina. The anhydrous aluminum chloride may be pre- 
pared by passing dry chlorine over the hot metal. The chloride 
volatilizes as fast as formed, thus constantly exposing a fresh 
surface of the metal to the gas. The chloride is made on a 
manufacturing scale by passing chlorine over a heated mixture 
of alumina and carbon : 



ALUMINUM. 299 

A1=0 /CI 

>0 + 3C + 6C1 = 2AK-C1 + 300. 
A1=0 \01 



The aluminum chloride distils and condenses in a suitable 
chamber. Common salt is often added to the mixture for the 
purpose of obtaining the double salt NaCl.AlCl 3 , which is also 
volatile. Sodium aluminum chloride is more permanent in 
air than aluminum chloride, and hence is better adapted for 
use in the manufacture of the metal. 

Aluminum Sodium Fluoride, AlF 3 .3NaF, occurs abundantly 
as the mineral cryolite in Greenland. It is extensively used 
in the manufacture of aluminum compounds and sodium 
hydroxide. 

Aluminum Oxide, Alumina, A1 2 3 or 0=Al-0-Al=0, oc- 
curs as the mineral corundum, of which the ruby, sapphire, 
oriental topaz, and oriental amethyst are varieties. Emery is 
granular gray or black corundum mixed with magnetite. 
Corundum ranks next to the diamond in hardness, hence its 
value in polishing and cutting. Aluminum oxide is prepared 
by igniting the hydroxide or aluminum salts of volatile acids. 
The artificial oxide dissolves with difficulty in acids, but may 
be rendered soluble by fusion with acid potassium sulphate or 
potassium hydroxide. 

Aluminum Hydroxides. — Ammonia produces in solutions of 
aluminum salts a bulky gelatinous precipitate, which consists 
of A1(0H) 3 together with some water, and which, on drying 
at common temperature, corresponds to A1(0H) 3 -f- H.,0. This 
loses water on heating above 300°, and aluminum oxyhydroxide, 
A10.0H, remains, which has the composition of the mineral 
diaspore. Precipitated aluminum hydroxide has the property 
of withdrawing dyestuifs and many salts from solutions. This 
property is applied in mordanting fabrics. 



300 THE THIKD GKOUP. 

Exp. 187. — Make a cochineal solution by digesting some powdered 
cochineal insects in hot water. Filter, and add a solution of alum, and 
then ammonia, and collect the precipitate on a filter. If sufficient alum 
is used the carmine color will be completely removed from the solution. 

Exp. 188. — Dip a piece of white cotton cloth into a dilute solution oi 
alum, then into water containing a little ammonia. Make a logwood 
solution by boiling some logwood chips in water, and place in it the 
cloth impregnated with aluminum hydroxide, and also a piece of white 
cotton cloth which has not been mordanted. Boil the solution some 
minutes and then wash and dry the pieces. 

A soluble aluminum h} r droxide has been obtained by decomposing 
aluminum acetate by long heating with water. The solution on evapo- 
ration yields a gelatinous residue which when dried at 100° corresponds 
to Al 2 0(OH) 4 . Graham dialyzed a solution of basic chloride and ob- 
tained a neutral solution of alumina, which after some days changed 
to a jelly. The addition of traces of salts, acids, and alkalies coagulate 
both these soluble hydroxides. 

Aluminates. — Aluminum hydroxide dissolves in solutions of 
the caustic alkalies with formation of compounds known as 
aluminates, in which aluminum acts as an acid radical. 

/O-Na 

Sodium Orthoaluminate, Na 3 A10 3 or Al(-0-Na. — Alumi- 

\0-Na 

num dissolves in a solution of sodium hydroxide, and for 
each atom of metal dissolved three atoms of hydrogen are 
liberated, thus : 

/O-Xa 
Al + 3Na-OH = Alf-O-Na + 3H. 
' M)-Na 

The orthoaluminate is said to separate when a solution of 
aluminum hydroxide in excess of sodium hydroxide is evapo- 
rated. 

Sodium Metaluminate, NaA10 2 or 0=Al-0-Na. — This com- 
pound is made by fusing cryolite with lime, and by melting to- 



ALUMINUM. 301 

getlier a native aluminum hydroxide with sodium hydroxide. 
The mass in either case is lixiviated with water ; the solution 
on evaporation yields the sodium aluminate as a white amor- 
phous powder. Sodium aluminate is used in mordanting 
fabrics,, and for the preparation of alumina. 

The solution of sodium aluminate is decomposed by carbon 
dioxide with separation of aluminum hydroxide : 

20=Al-0-Na + C0 2 + 3H 2 = 2A1^0H + CO<£^* 
Calcium hydroxide precipitates calcium aluminate : 
20=Al-0-Na + Ca<Qg = o=AM) >0a + 2Na-OH. 

The strontium and barium aluminates are soluble in water. 
The potassium compound 20 = Al-0-K -|- 3H 2 has been 
obtained in crystals. 

Exp. 189. — Add sodium hydroxide cautiously to an alum solution 
until the precipitate which at first forms dissolves. Pass into the so- 
lution obtained carbon dioxide, or treat a solution of commercial sodium 
aluminate with carbon dioxide. Express the reactions by equations. 

so »<o)ai 

Aluminum Sulphate, A1 2 (S0 4 ) 3 or S0,.<q .—This salt 

so 5 <° ) A1 

is prepared by dissolving aluminum hydroxide in sulphuric 
acid, and also from clay. The clay, as free as possible from 
iron, is roasted to render the iron present less soluble, and the 
aluminum silicate more easily decomposed by the acid. It is 
then heated with sulphuric acid, and the solution, after the 
silica and umdeeom posed clay have subsided, is drawn otV and 



302 THE THIRD GROUP. 

evaporated until it solidifies on cooling. This crude sulphate 
is used by dyers as a mordant. It is also used for " weighting- 
paper. " 

Aluminum sulphate crystallizes with 16 molecules of 
water in pearly scales, but if ferric sulphate is present it is 
hygroscopic and contains rather more than 18 molecules of 
water. Freshly precipitated aluminum hydroxide dissolves 
in a solution of aluminum sulphate, with the formation doubt- 
less of basic sulphate. 

Alums. — Aluminum sulphate forms with alkali sulphates 
double salts which crystallize in octahedrons, usually exhibit- 
ing cubic faces. The name alum was formerly applied only 
to the potassium aluminum sulphate, but the term is now 
used to designate a group of analogous bodies, in which alumi- 
num is replaced by iron, chromium, and other metals, and 
the potassium by other alkali metals, or by silver, thallium, 
or ammonium. 

Potassium Alum, A1 2 (S0 4 ) 3 .K 2 S0 4 + 24H 2 0.— If a solution 
of aluminum sulphate is mixed with a solution of potassium 
sulphate in sufficient quantity to form the double salt, no 
change in temperature will occur — an indication that there has 
been no combination. The solution of the two salts, however, 
yields on evaporation crystals of alum. 

At 0° 100 parts of water dissolve 3.9, at 20° 15, and at 100° 
357.5 parts of A1 2 (S0 4 ) 3 .K 2 S0 4 + 24H 2 0. Alum melts at 92° 
in its water of crystallization, most of which it loses slowly at 
100°, and the last portions on gentle ignition. The porous 
mass thus obtained, known as burnt alum, dissolves slowly in 
water. If potassium hydroxide is slowly added to potash 
alum, so long as the precipitate redissolves readily on stir- 
ring, a neutral solution is obtained, which is known in 
the arts as neutral alum. Solutions of common alum, to 
which a small quantity of neutral alum has been added, yield 



GALLIUM. 303 

on spontaneous evaporation cubic crystals of the same compo- 
sition as octahedral alum. 

Ammonium Alum, A1 2 (S0 4 ) 3 .(NH 4 ) 2 S0 4 + 24H 2 0, has been 
manufactured at times when ammonium sulphate has been 
cheaper than potassium sulphate. It is somewhat more 
soluble in water than potassium alum, which it closely re- 
sembles. Commercial alum sometimes contains both potas- 
sium and ammonium. Alum is valuable as a source of alumi- 
num hydroxide free from iron, but it is largely replaced for 
use in dyeing by aluminum sulphate and sodium aluminate. 

Exp. 190 —Take 20 grams crystallized aluminum sulphate and the 
weight of ammonium sulphate required to form ammonium alum. 
Dissolve the two salts together in 100° cc. of boiling water, filter the 
solution into a porcelain dish, and let stand for several days to crys- 
tallize. Test one of the crystals for ammonia and also for aluminum. 



Gallium^ GU. 

Atomic Weight, 70. Density, 5.9. 

Gallium was discovered by means of the spectroscope in 1875 by 
Lecoq de Boisbaudran in zinc blende from the mine of Pierrefitte in 
the Pyrenees. The blende of Bensberg has since proved to be a better 
source of the metal, and from 43,000 kilos, of the ore 62 grams of gal- 
lium were obtained. The metal is separated by electrolyzing a solution 
of a salt of it in potassium hydroxide. 

Gallium melts at 30° to a brilliant white liquid, which rnay be cooled 
10°-15° below its fusing point without solidifying. In the solid state 
the metal has a bluish tinge, is hard, and but slightly malleable. It 
does not lose its lustre in air or in water free from air, but slowly tar- 
nishes in ordinary water. It oxidizes superficially at a red heat in the air. 
and does not volatilize. Cold nitric acid is without action on the metal, 
but the hot acid dissolves it slowly. Hydrochloric acid dissolves it 
readily and potassium hydroxide slowly with evolution of hydrogen. 



304 THE THIRD GROUP. 

Gallium Dichloride, GaCl 2 . — Gallium is readily attacked by chlorine 
with formation of gallium dichloride if the metal is in excess and the 
current of chlorine moderate. With an excess of chlorine the trichlor- 
ide, GaCl 3 , is formed. The latter is the more volatile of the two com- 
pounds. Gallium dichloride is decomposed by water with evolution 
of hydrogen and formation of the trichloride. 

Gallium Trichloride, GaCl 3 , fumes strongty in the air and quickly deli- 
quesces. It fuses at 75°. 5 and boils at 215°-220\ Its gas density at 
273° is 171.8, corresponding to gas molecules of Ga 2 Cl 6 . At high tem- 
peratures the density found equals 95.3, showing dissociation into mole- 
cules of GaCl 3 . At still higher temperatures it decomposes. 

Gallium Oxide, Ga 2 3 , is obtained by igniting the nitrate. Ammonia 
added to a solution of a gallium salt precipitates gallium oxide (hydrox- 
ide?), and the precipitate is more soluble in ammonia than aluminum 
hydroxide. Potassium hydroxide also precipitates the oxide, which 
readily dissolves in an excess of the precipitant. From this alkaline 
solution carbonic acid again precipitates the oxide 

Gallium Nitrate, Ga(N0 3 ) 3 , is obtained by evaporating at 40° a solution 
of gallium in nitric acid. 

Gallium Sulphate, Ga 2 (S0 4 )3, is very soluble in water, and forms with 
ammonium sulphate gallium-ammonium alum, which crystallizes in 
octahedrons or octahedrons with cubic facets. 



Indium j In. 

Atomic Weight, 113.7. Density, 7.4. 

Indium occurs in small quantities in the zinc blende of various locali- 
ties It was discovered by means of the spectroscope in 1863 by Reich 
and Richter in the blende of Freiberg, Saxony, which contains about 
0.1 per cent of indium. 

Indium is a silver-white non-crystalline metal, softer than lead. It 
melts at 176° and oxidizes at a higher temperature. At a red heat it 
burns with a blue flame to the oxide. It is less volatile than zinc and 
cadmium, and is not corroded by moist air or boiling water. It dis- 
solves slowly in dilute hydrochloric and sulphuric acids, and readily in 
nitric acid. 

Indium Chloride, InCl 3 , is prepared by burning the metal in chlorine. 
It is very deliquescent. 



THALLIUM. 305 

The observed gas density of indium chloride is 113.6; theory re- 
quires 109.9 for InCl 3 . 

Indium Oxide, ln 2 3 , is obtained by iguitiou of the nitrate, and also 
by burning the metal. It is a pale-yellow powder, which turns brown 
on heating. 

Indium Hydroxide, In(0H) 3 . — Ammonia produces in solutions of in- 
dium salts a bulky precipitate similar to aluminum hydroxide. It has 
the composition of In(OH) 3 after drying at 100°. It is soluble in acids. 
Potassium and sodium hydroxides dissolve it, but the solution soon be- 
comes turbid owing to separation of indium hydroxide. 

Indium Sulphate, In 2 (S0 4 )3, is obtained by dissolving the oxide in sul- 
phuric acid. With ammonium sulphate it forms the alum In 2 (S0 4 ) 3 . 
(NH 4 ) 2 S0 4 + 24H 3 0. 



Thallium, Tl. 

Atomic Weight, 204. Density, 11.8. 

Thallium is widely distributed, although found only in 
small quantities. It occurs in copper and iron pyrites and 
other sulphides of many localities, and has been found in 
some spring waters. Thallium is best obtained from dust 
which collects in the flues to the chambers in which thallifer- 
ous pyrites is burned in the manufacture of sulphuric acid. 
The dust seldom contains more than a fraction of one per 
cent of the metal. 

Thallium was discovered by Orookes in a selenif erous deposit 
from a sulphuric-acid factory in the Harz. From this deposit 
selenium was prepared. This left on distilling a residue that 
gave in the spectroscope a single green line hitherto unknown. 

Metallic thallium is separated from its solutions by zinc or 
by electrolysis, and also by fusing the iodide with potassium 
cyanide. It is not quite as white as silver, is softer than lead, 
melts at 290°, and may be distilled at a white heat in hydro- 
gen. The bright metal tarnishes quickly in air, becoming 
covered with a thin film of thallous oxide, which protects it 
from further oxidation. 



306 THE THIRD GROUP. 

Thallium forms two series of compounds : the thallous, in 
which it is univalent ; and the thallic compounds, in which it 
is trivalent. In thallous compounds thallium presents close 
analogy in properties to the alkali metals and to silver. 



Thallous Compounds. 

Thallous Chloride, T1C1, forms a white precipitate when 
hydrochloric acid is added to a solution of a thallous salt. It 
dissolves in 63 parts of water at 100° and in 504 parts at 0°. 
It crystallizes in cubes from hot solutions. Two determina- 
tions of its gas density gave 106.8 and 126 ; T1C1 requires 119. 

Thallous Iodide, Til, is very insoluble in water, and thallium 
may be separated as iodide from dilute solutions by addition 
of potassium iodide. 

Thallous Oxide, T1 2 or T1-0-T1, is obtained by heating 
thallous hydroxide to 100° out of contact with air. It is a 
black powder, which absorbs moisture from air, and dissolves 
in water with formation of the hydroxide. 

Thallous Hydroxide, T1-0H, is best prepared by decompos- 
ing a solution of thallous sulphate with barium hydroxide. 
The solution filtered from the barium sulphate yields on evapo- 
ration yellow crystals, having the composition TIOH + H 2 0. 
The solution of thallous hydroxide has a strong alkaline re- 
action, and is similar to potassium hydroxide in its deport- 
ment towards solutions of metallic salts. 

Thallous Nitrate, T1N0 3 or NO -0-T1, is formed together with 
a little thallic nitrate when thallium is dissolved in nitric acid. 
It is a soluble salt, which melts without decomposition at 205°. 



THALLIC COMPOUNDS. 307 

O-Tl 

Thallous Sulphate, T1 2 S0 4 or S0 2 <q r^.— This salt is ob- 
tained by dissolving the metal or thallous hydroxide in sul- 
phuric acid. It crystallizes in rhombic prisms, isomorphous 
with potassium sulphate. With aluminum sulphate it forms 
the alum A1 2 (S0 4 ) 3 .T1 2 S0 4 + 24H 2 0. Hydrogen thallous sul- 
phate, HT1S0 4 , has been prepared. 

Thallous Sulphide, T1 2 S, separates as a black precipitate 
when an alkaline solution of a thallous salt is treated with 
hydrogen sulphide. 



Thallic Compounds. 

Thallic Oxide, T1 2 3 , is formed when thallium burns in 
oxygen and when thallic hydroxide is heated to 100°-115°. 
It gives off oxygen at a strong red heat, and is insoluble in 
water and in caustic alkalies. It dissolves in warm sulphuric 
acid with evolution of oxygen. 

Thallic Hydroxide. — Ammonia and potassium hydroxide 
produce in a solution of thallic chloride a precipitate which 
on drying has the composition TIO.OH (= = Tl-OH). 



/CI 
Thallic Chloride, T1C1 3 or Tl— CI.— This compound is formed 

when thallic hydroxide is dissolved in cold hydrochloric acid, 
or by treating thallium or thallous chloride under water with 
chlorine. The solution on evaporation in vacuum yields hy- 
drous crystals, which become anhydrous at 50°-C0°, and at 
100° lose chlorine. 



308 THE THIKD GROUP. 

so *<o)ti 

Thallic Sulphate, 11,(80,) , + 7H,0 or SO^ , + 7H 2 0. 

so,<^ 

— This salt separates in crystals when a solution of thallic hy- 
droxide in sulphuric acid is evaporated. The salt is decom- 
posed by water with separation of thallic hydroxide. 

N0-0 X 

Thallic Nitrate, T1(N0 3 ) 3 or NO^-O-^Tl, separates in hy- 

NO-CK 

drous deliquescent crystals from a solution of the hydroxide 
in nitric acid of density 1.40. They decompose on heating 
with water. 



Scandium, Sc. 

Atomic Weight, 44. 

This element was discovered by Mlsou in 1 879 in his attempt to 
purify ytterbium oxide, which was obtained mostly from the minerals 
gadolinite and euxenite. After repeating a number of times a process 
for separating the oxides of erbium and ytterbium, he finally obtained 
a small residue of earthy oxide, which he found to possess a lower 
combining weight than the oxides of erbium and ytterbium. This 
residue was submitted to Thalen, who found in addition to the spectra 
of erbium and ytterbium a number of lines of the new element. 

Scandium Oxide, Sc 2 3 , is a white, light, infusible powder, in appear- 
ance much like magnesia. It is obtained by the ignition of the hydrox- 
ide, nitrate, or sulphate. It dissolves in boiling concentrated nitric or 
hydrochloric acid. Addition of ammonia to a solution of a scandium 
salt precipitates scandium hydroxide, which is insoluble in caustic 
alkalies. 

Scandium Nitrate crystallizes from solutions of the salt which have 
been evaporated to syrupy consistence. It gives off nitric acid on 
heating, and a basic salt is formed which is completely soluble in water. 
If the decomposition is carried so far that only a little nitric acid 
remains, the residue renders water milky, and the water does not 



YTTRIUM — LAKTHAHUM. 309 

become clear even on long standing. Nilson regards this as character- 
istic of scandium. 

Scandium Sulphate, Sc 2 (S0 4 ) 3 , is prepared by heating the nitrate with 
sulphuric acid. It crystallizes from aqueous solutions with six molecules 
of water. The scandium potassium sulphate, Sc 2 K 6 (S0 4 ) 6 , is insoluble 
in a saturated solution of potassium sulphate. 



Yttrium, Y. 

Atomic Weight, 89. 

Yttrium has not been separated in the pure state. 

Yttrium Oxide, Y 2 3 , is a yellowish-white powder prepared by igniting 
the oxalate or hydroxide. It does not unite directly with water, but 
the hydroxide is obtained as a gelatinous precipitate when a caustic 
alkali is added to a solution of an yttrium salt. The hydroxide decom- 
poses ammonium salts, and like the alkalies combines directly with 
carbon dioxide. 

Yttrium Chloride, YC1 3 . — Yttrium oxide dissolves slowly in hydro- 
chloric acid, and the solution on evaporation yields crystals of the com- 
pound YC1 3 -|- 6H 2 0, which becomes anhydrous when ignited with 
ammonium chloride. 

Yttrium Nitrate, Y(N0 3 ) 3 + 6H 2 0, forms needle-shaped crystals which 
do not change in air. 

Yttrium Sulphate, Y 2 (S0 4 ) 3 + 8H 2 — This salt is more soluble in cold 
than in warm water. 100 parts of water at 15°. 5 dissolve 15.2 parts of 
the anhydrous salt, but on warming a portion of the salt crystallizes 
out. 

Alums do not appear to have been obtained with yttrium sulphate, 
but double salts have been prepared with potassium and ammonium 
sulphates. 



Lanthanum, La. 

Atomic Weight, 139. Density, 6.1. 

Lanthanum has been found in a few minerals associated with cerium 
and didymium. The metal was obtained by Hildebrand and Morton 



310 THE THIRD GROUP. 

by electrolyzing lanthanum chloride. It is readily acted on by strong 
nitric acid, and tarnishes quickly in dry air. It is malleable, but can- 
not be drawn into wire. The finely divided metal burns brilliantly in 
air. It decomposes water with evolution of hydrogen and formation 
of lanthanum hydroxide. 

Lanthanum Oxide, La 2 3 , is obtained by igniting the nitrate. It be- 
comes warm when mixed with water, with formation of lanthanum 
hydroxide, La(OH) 3 . The hydroxide reacts alkaline, absorbs carbon 
dioxide from the air, and is precipitated from solutions by caustic 
alkalies. 

Lanthanum Chloride, LaCl 3 , is a soluble salt which crystallizes with 
seven molecules of water. 

Lanthanum Nitrate, La(N0 3 ) 3 + 6H 2 0, is an easily soluble salt. 

Lanthanum Sulphate, La 2 (S0 4 ) 3 -f- 9H 2 0, is much more soluble in cold 
than hot water. 

Lanthanum Potassium Sulphate is precipitated when solutions of the 
two sulphates are mixed. It is completely insoluble in a solution of 
potassium sulphate. 

Lanthanum Ammonium Sulphate, La 2 (NH 4 ) 2 (S0 4 ) 4 + 8H 2 0, crystallizes 
in flattened striated prisms, soluble in water and permanent in air. 



Erbium, Er. 

Atomic Weight, 166. 

This element occurs very sparingly associated with ytterbium and 
scandium. The metal is obtained as a gray powder from the chloride 
of erbium and sodium by electrolysis, and by reduction with sodium. 
It decomposes water with evolution of hydrogen. 

Erbium Chloride, ErCl 3 , is obtained by heating out of contact with air 
the residue from the evaporation of a solution of the oxide in hydro- 
chloric acid, to which ammonium chloride has been added. It crystal- 
lizes from solution with six molecules of water. 

Erbium Oxide, Er 2 3 , is a pale rose-colored powder obtained by ignition 
of the nitrate. 

Erbium Nitrate, Er(N0 3 ) 3 -|- 6H 2 0, separates in large crystals when a 
solution of the oxide in nitric acid is evaporated. 

Erbium Sulphate, Er 2 (S0 4 ) 3 , crystallizes with eight molecules of water, 



SUMMARY OF THE THIRD GROUP. 311 

which are given off above 100°. The anhydrous salt is more readily 
soluble in water than the hydrous salt. With potassium sulphate it 
forms the double salt K 2 S0 4 .Er 2 (S0 4 )3 + 4H 2 0. 



Ytterbium, Yb. 

Atomic Weight, 173. 

Ytterbium has not been obtained in the metallic state. It is a very 
rare element, found in only a few minerals. Euxenite thus far has been 
the best source of it. 

Ytterbium Oxide, Yb 2 3 , is a white, heavy, infusible powder, which 
dissolves slowly in cold and easily in boiling dilute acids. 

Ammonia precipitates from ytterbium solutions a hydroxide which 
absorbs carbon dioxide from the air. The hydroxide is easily soluble 
in acids, and on ignition leaves the oxide. 

Ytterbium Nitrate changes on heating to a basic salt readily soluble in 
water. 

Ytterbium Sulphate, Yb 2 (S0 4 ) 3 , is obtained by heating the nitrate with 
sulphuric acid until the excess of acid is expelled. A solution of the 
anhydrous salt in water yields on evaporation crystals of Yb 2 (S0 4 )3 + 
8H 2 0, which are permanent in air. 



Summary of the Third Group. 

Boron, with the lowest atomic weight in the group, stands 
apart from the other members. It is non-metallic and acidic 
in character, and does not exhibit basic properties. Alumi- 
num, with the next lowest atomic weight in the group, is both 
basic and acidic, showing the latter property in the alu mi- 
nates. These compounds are analogous in composition to the 
ortho- and meta-borates. The ortho- and meta-aluminates 
are decomposed by carbon dioxide with separation of alumi- 
num hydroxide, whereas the corresponding alkali borates are 



312 THE THIRD GROUP. 

converted by carbon dioxide into tetraborates. Gallium, 
indium, and thallium possess increasing basic properties in 
accord with their increasing atomic weights. Thallium, with 
the highest atomic weight, stands apart from the other mem- 
bers. The thallous halides resemble in insolubility the halides 
of silver and mercury, and thallous hydroxide is a strong base, 
which separates metallic hydroxides from salts. 

The rare elements scandium, yttrium, lanthanum, erbium, 
and ytterbium are not as well understood as the other members 
of the group. The oxides of the group, thallium oxides and 
boric oxide excepted, are known as earthy oxides. 

The student should tabulate the compounds of the group. 



THE FOURTH GROUP. 

The members of this group are carbon, silicon, tin, and 
lead, and the rare elements titanium, germanium, zirconium, 
cerium, and thorium. 



Carbon, C 

Atomic Weight, 12. 



Carbon is a constituent of animal and plant matter, and of 
all organic compounds. It occurs in the air as carbon dioxide, 
and in the earth in rock masses of carbonates, and in the 
various forms of coal. 

It exists in different allotropic modifications, which are clas- 
sified as diamond, graphite, and amorphous carbon. Under 
the last are included gas carbon, anthracite coal, coke, 
charcoal, and lamp-black. These varieties differ in color, 
form, density, conductivity for heat and electricity, and in 
their behavior towards oxygen. The diamond, graphite, and 
gas carbon burn with difficulty, whilst charcoal and coke are 
easily combustible. All of the forms of carbon burn with 
oxygen to carbon dioxide. Charcoal is a poor conductor 
of electricity; graphite, coke, and gas carbon arc good con- 
ductors, but the diamond does not conduct electricity. Graph- 
ite, and the denser forms of amorphous carbon, such as coke 
and anthracite, conduct heat better than charcoal. Carbon 
is infusible and non-volatile except in the electric arc, and 
is insoluble in all ordinary solvents, and resists the combined 
corrosive action of air and moisture. 



314 



THE FOURTH GROUP. 




Diamond owes its value to its hardness and the bril- 
with which it reflects light. It is found crystallized 
in octahedral, dodecahedral, and more complex 
forms, often with curved edges. It varies in 
color from colorless, yellowish, green, red, blue 
to black. The density of pure specimens is 3.5, 
that of the black diamond is 3 or less. The dia- 
mond burns when heated intensely in air or 
oxygen to carbon dioxide, and leaves a small 
amount of incombustible ash. 

Exp. 191. — Enclose a small fragment of a diamond in 
a coil (Fig. 86) of small platinum wire connected with 
copper wire passing through a rubber stopper in the 
cylinder containing oxygen and some clear lime water. 
On heating the platinum with a battery current the 
Fig. 88. diamond will take fire and burn brilliantly, and the 
lime water will react for carbon dioxide. 

The diamond remains unchanged at a white heat out of 
contact with air, but at the intense heat of the electric arc it 
swells and changes to a coke-like mass. 

As the diamond is the hardest of known substances, it can 
only be cut and polished by diamond dust. Only the trans- 
parent and more perfect stones are valuable as gems, which 
are cut so as to best reflect light. All the light which strikes 
the back planes at an angle greater than 24° 13' is reflected. 
The weight of diamonds is commonly stated in carats, the 
carat being equal to 3.17 grains and 0.2054 gram. 

Only the natural curved edges of diamonds answer for 
cutting glass. The cutting diamond simply makes a small 
crack in the glass, which determines the line of fracture. 
Writing diamonds are made of small splinters, which scratch 
but do not crack glass. Diamond dust is used in cutting 
and polishing precious stones. The black Brazilian diamonds 
have within a few years been employed in drilling rock, and 
consequently the price of them is much higher than formerly. 



CAEBOK. 315 

Graphite, also called plumbago and black-lead, occurs in 
foliated and in compact and granular masses, more rarely in 
hexagonal prisms. It is formed artificially in several ways. 
Molten cast iron dissolves more carbon than it can hold in 
combination on cooling, and the carbon partly separates as 
scales of graphite when the iron solidifies. Graphite is also 
produced when the concentrated black-ash liquors of soda 
manufacture are oxidized with nitre. Graphite has a brilliant 
metallic lustre, varying from an iron black to a steel gray. 
Its density is 2.5. It feels soft and greasy, and leaves a black 
mark when rubbed. Natural graphite contains earthy matters 
as impurities, and usually 0.5 to 1.3 per cent of hydrogen. It 
is purified by treatment with aqua regia and hydrofluoric 
acid, or with a mixture of sulphuric and nitric acids ; also by 
heating with potassium chlorate and sulphuric acid. In the 
last process the residue is washed with water, and the prod- 
uct, which contains carbon, hydrogen, oxygen, and sulphuric 
acid, is dried and then heated, when it gives off gas and leaves 
pure graphite in the form of a fine powder. Graphite is less 
combustible "than the diamond, and burns slowly and only at 
high temperatures in air and oxygen. 

Exp. 192. — Cut the wood away from the lead of a pencil for two 
centimeters. Heat the end of the lead to redness in the lamp flame, 
and observe whether it burns or not. 

The leads for leadpencils were formerly cut from blocks 
of compact graphite. In the improved process powdered 
graphite is mixed with fine clay, and the plastic mass is forced 
by great pressure through a hole, and thus formed into the 
shape required for pencils. Black-lead crucibles are made of 
a mixture of graphite and fire-clay. They are better conduc- 
tors of heat, are more refractory, and less liable to crack 
than clay crucibles. They are much used in melting steel. 
Graphite is used as a lubricant and as a polishing powder, and 
also for facing sand moulds in iron foundries. It is the chief 



316 THE FOURTH GBOUP. 

ingredient of stove-blacking, and not only gives a good lustre 
but also prevents the iron from rusting. In the electrotype 
process the moulds are coated with powdered graphite, which 
serves as a conductor of electricity. 

Gas Carbon collects on the upper part of the interior of the 
retorts in which coal is distilled in the manufacture of gas. 
The hydrocarbon vapors and gases are partially decomposed 
by the intense heat with separation of carbon. A similar re- 
sult is obtained by passing ethylene gas, C 2 H 4 , one of the con- 
stituents of coal gas, through a white-hot porcelain tube. Part 
of the carbon deposits, and free hydrogen and complex hydro- 
carbons are formed. Gas carbon has a gray metallic lustre, 
and is very hard. Its density varies from 1.7 to 2.5. It is 
used for the carbon plates of the Bunsen battery, and for the 
carbon poles of the arc electric lamps. 

Coke. — Certain varieties of bituminous coal when heated 
swell, become pasty, and after gases are no longer evolved, a 
porous mass remains, called coke. It is hard, has a grayish 
metallic lustre, and is a valuable fuel. It is formed in the 
retort in the manufacture of coal gas, and is made in enor- 
mous quantities by heating coal in suitable ovens. Coke con- 
tains besides carbon a little hydrogen, oxygen, and nitrogen, 
and the ash which was contained in the coal. 

Lamp-black. — When a cold body is held in a luminous gas 
or lamp flame soot or lamp-black deposits. This is owing to 
the fact that the hydrogen burns first, leaving the carbon on 
the cold surface. Substances rich in carbon, such as kerosene 
oil, burn with smoky flames, and the amount of carbon sepa- 
rated may be increased by diminishing the supply of air. 
Lamp-black is manufactured by burning tar, rosin, turpen- 
tine, or petroleum with a small supply of air. The products 
of combustion pass through large chambers hung with coarse 



CAKBON. 317 

sacking to collect the lamp-black. It is the basis cf printing 
ink, and is used as a black pigment for various purposes. It 
contains hydrocarbons which may be partly expelled by igni- 
tion, but in order to remove the last traces of hydrogen it 
must be heated in chlorine, which combines with the hydro- 
gen, but does not act upon the carbon. 

Exp. 193. — Place a burning turpentine lamp under a bell- jar on a 
plate. Much lamp-black will separate. 

Charcoal. — When pure sugar is heated in a platinum dish a 
porous glistening black residue remains, which contains a little 
oxygen and hydrogen even after intensely heating. By ignit- 
ing this charcoal in pure chlorine pure carbon is obtained. It 
is tasteless, insoluble, conducts electricity, and has a density of 
1.57. Sugar charcoal is useful in the laboratory when carbon 
free from ash is required. 

Wood charcoal is made by covering a pile of wood with 
charcoal dust and moist earth. The wood is set on fire in suit- 
able openings and allowed to burn slowly for some time ; then 
the openings are closed, and the pile is left until cold. A 
better yield of charcoal is obtained by subjecting wood to what 
is known as dry distillation in retorts or brick chambers, a 
process which admits of saving the liquid products of the dis- 
tillation. Wood when dried at 150° contains about half its 
weight of carbon, the other half consisting of hydrogen and 
oxygen in nearly the proportion to form water, and a little 
nitrogen and ash. When the wood is charred, part of the 
carbon passes off with the volatile products, and the coal re- 
maining differs in composition according to the temperature 
employed, as the following results of Violette's experiments 
with wood dried at 150° show : 



318 



THE FOURTH GROUP. 



Percentage 
Temperature. yield of 

C harcoal. 

150° 

280°., 36.2 

350° 29.7 

432° 18.9 

1032° 18.7 

1160° 18.4 

1300° 17.5 

1500° 17.3 

Over 1500° 15.0 



Carbon Hydrogen. 



47.5 
71.6 
76.6 
81.6 
81.9 
83.3 
90.8 
94.5 
96.5 



6.1 

4.7 
41 
1.9 
2.3 
1.7 
1:6 
0.7 
0.6 



Oxygen 
and Nitrogen. 

46 3 
22.1 

18.4 
15.2 
14.1 
13.8 

6.5 

3.8 

0.9 



Ash. 

0.08 

0.57 

0.6 

1.2 

1.6 

1.2 

1.1 

0.7 

1.9 



The table shows that charcoal made at low temperatures 
retains considerable hydrogen and ox}^gen, and that the 
amount of the product is diminished with increasing tem- 
perature of charring. Slow charring yields more charcoal 
than rapid, and dry wood more than wet. Heavy dense 
woods yield a denser and less porous coal than light soft woods. 
Charcoal made at 300° is brown and soft, and takes fire at 
about 380° ; while that made at 1000°-1500° is black, hard, 
and brittle, and ignites at about 700°. 

All solids possess the property of condensing gases on their 
surfaces at ordinary temperatures. Wood charcoal, owing to 
its porous structure, has an enormous extent of surface within 
its pores, and has the property in a marked degree of absorb- 
ing or condensing gases. The more porous the charcoal the 
greater the quantity of gas it will absorb. Hunter found that 
one volume of charcoal absorbed the following quantities of 
gases (reduced to 0° and 760 mm. pressure): 





Volumes 


Ammonia, . . . 


. 171.7 


Cyanogen, . . . 


. 107.5 


Mtrous oxide, . . 


. 86.3 


Ethene, .... 


. 74.4 


Nitric oxide, . , 


. 70.5 



Volumes. 

Carbon dioxide, . . .67.7 
Carbon monoxide, . .21.2 

Oxygen, 17.9 

Nitrogen, 15.2 

Hydrogen, . , , , 4.4 



CATLBOST. 319 

The results show that the more readily condensible gases 
are absorbed in the greatest quantity. Gases contained in 
charcoal may be expelled by heat, and are mostly given off in 
a vacuum. Freshly ignited charcoal absorbs the gases of the 
atmosphere and aqueous vapor, and if the air is damp the 
charcoal takes up considerable water. The well-known prop- 
erty that charcoal possesses of removing noxious gases is due 
not only to their absorption within the pores, but to their 
oxidation to carbon dioxide and water by oxygen also ab- 
sorbed. These products diffuse into the air almost as soon as 
formed, and the charcoal continues to absorb the noxious gas 
and oxygen, and bring about their combination. Thus it is 
that charcoal has the property of burning up the foul gases of 
the decay of bodies greater in weight than the charcoal used. 
It should be understood that charcoal, in the action just de- 
scribed, is not an antiseptic, but is simply a self-acting cre- 
matory, in which baneful products of decay are oxidized to the 
harmless products carbon dioxide and water. Charcoal re- 
tains this property for a long time, and when it has become 
impaired it can be restored by simply heating to redness. 

Exp, 194.— Fill a tube over mercury with ammonia gas. Heat a 
cylindrical piece of charcoal to redness in a platinum crucible to expel 
the gases from it ; then take the glowing coal with a pair of crucible 
tongs and put it into the glass tube without lifting the latter above the 
surface of the mercury. The gas will be rapidly absorbed, aud the 
mercury will soon rise to the top of the tube. On removing the char- 
coal from the tube it will be evident from the odor that the ammonia 
which was absorbed is rapidly given off. 

Exp. 195.— Place a small crucible filled with freshly ignited and 
nearly cold charcoal powder in a jar of hydrogen sulphide, then put it 
into a jar of oxygen. The rapid oxidation of the hydrogen sulphide 
will ignite the charcoal. 

Exp. 196.— Hold a piece of charcoal under hot water in a test-tube 
by means of a glass tube or rod. Air will bubble from the charcoal, 
and if the latter is held under hot water long enough it will not tloat. 

Exp. 197.— Place pieces of meat in three glass cylinders, and above 
the meat put a piece of wire gauze and a thin layer of cotton. Cover 



320 THE FOURTH GROUP. 






the cotton in one cylinder with an inch of powdered charcoal. In the 
second cylinder place the same depth of dry fine earth, and in the third 
put dry sand. The meat will in a few days show marked indications of 
decay : that covered with sand will not have changed more than the 
others, but an odor will come from it ; while no odor or a little ammo- 
nia will be noticed in the cylinders containing the charcoal and earth. 

Animal Charcoal. — Wood charcoal has the property of with- 
drawing certain matters from solution and absorbing them in 
its pores. Animal charcoal possesses this power to a greater 
extent. It is made by calcining bones and other animal mat- 
ter out of contact with air, and differs from wood charcoal in 
that it contains nitrogen. Bone-black is a black porous mass, 
consisting chiefly of calcium phosphate and carbon, obtained 
by heating bones. It is largely used in sugar-refining to re- 
move coloring matter and lime salts from solutions of sugar. 
Bone-black is not adapted for decolorizing acid solutions which 
dissolve calcium phosphate, but blood charcoal is not open to 
this objection. It is made by evaporating and calcining a 
mixture of blood and potassium carbonate. The residue is 
exhausted with boiling Avater and hydrochloric acid, to remove 
the soluble portions, including the potassium salt which was 
added to make the product very porous. Aniinal charcoal is 
used in the laboratory to purify organic preparations, and in 
some cases to separate compounds from solutions. The fol- 
lowing experiments illustrate some of the properties and uses 
of it: 

Exp 198.— Treat a highly diluted solution of acid quinine sulphate 
with boue-black. After a time the solution will not taste of quinine. 

Exp. 199. — Decolorize with bone-black a very dilute solution of the 
coloring matter obtained by digesting logwood chips with water. 

Mineral Coal is mainly of vegetable origin. The coal beds 
are the result of the slow decomposition under water and 
earth of plants, the change consisting of a loss of carbon di- 
oxide, water, and marsh gas, and the formation of products 



COMPOUNDS OF CARBON". 321 

richer in carbon than the woody fibre of the plants. Coals 
are classified as non-flaming and flaming, according as they 
burn without or with a luminous flame. Anthracite, a non- 
flaming coal, burns with a pale-blue flame. It consists mainly 
of carbon, a little hydrogen, oxygen, nitrogen, and earthy 
matter or ash, and it yields from 4 to 7 per cent of volatile 
matter. The flaming coals contain more hydrogen and oxy- 
gen than anthracite. The varieties are: bituminous coal, 
which yields from 20 to 40 per cent of volatile matter ; cannel 
or candle coal, so named because a small fragment burns 
readily like a candle, with 50 to 60 per cent of volatile ingre- 
dients ; and brown coal, which is richer in oxygen than the 
other varieties. Peat, which is forming at the present time in 
bogs, more nearly approaches wood in composition. 

The following are some analyses of coal and peat, and for 
the sake of comparison the composition of wood is also given : 







Hydro- 




Nitro- 








Carbon. 


gen. 


Oxygen. 


gen. 


Sulphur, 


, Ash. 


Wood, excluding ash, 


. 50.00 


6.00 


44.00 








Peat, 


. 59.70 


5.70 


33.04 


1.56 






Brown Coal, Bovey, . . 


. 66.31 


5.63 


22.86 


0.57 


2 36 


2.27 


Cannel Coal, Wigan, . . 


. 80.07 


5.53 


8.10 


2.12 


1.50 


2.70 


Non-coking, Brial Hill, 0. 


. 78.94 


5.92 


11.50 


1.58 


0.56 


1.45 


Non-coking, Indiana, 


. 82.70 


4.77 


9.39 


1.62 


0.45 


1.07 


Coking Coal, Kentucky, 


. 74.45 


4.93 


13.08 


1.03 


0.91 


5.00 


Anthracite, Pennsylvania, 


. 90.45 


2.43 


2.45 






4.67 


<< (< 


. 92.59 


2.63 


1.61 


0.92 




2.25 



Compounds of Carbon. 

There are an enormous number of carbon compounds. With 
few exceptions they contain hydrogen and oxygon, and less 
frequently nitrogen. Many carbon compounds are formed by 
animal and plant growth, and from them many others are 



322 THE FOURTH GROUP. 

prepared. Animals and plants develop from germs, and ex- 
hibit an organized structure: hence the carbon compounds 
derived from them have been termed organic. The study of 
the carbon compounds constitutes a branch of chemical science 
known as Organic Chemistry. Formerly compounds of carbon 
and hydrogen were only obtained from animal and vegetable 
substances; now a number of them can be formed syntheti- 
cally, showing that there is no sharp distinction to be made 
between inorganic and organic compounds. 

Carbon is tetravalent, and rarely exhibits a lower valence. 
Its oxides will be treated of first, then their derivatives; a few 
of the compounds of carbon and nitrogen will be described, 
and also a few of the simpler organic compounds. 

Isomerism. — Compounds having the same composition and 
molecular weight, but differing in physical properties and 
chemical deportment, are said to be isomeric. The differences 
which isomers present are due to the different relations exist- 
ing between the atoms in the molecule. While the positions 
of the atoms in a molecule are unknown, the relations of the 
atoms in many kinds of molecules are fixed with some degree 
of certainty, and are expressed by structural formulas. The 
existence of isomers renders some way of expressing their differ- 
ences in constitution a necessity. For example, there are two 
isomeric compounds having the empirical formula C 2 H 6 0. 
One, methyl oxide, is a gas at common temperature, and the 
other is ethyl hydroxide or common alcohol. Their differences 
in chemical character are-expressed by the formulas 



H H 


HH 


H-C-O-C-H; 


H-C-C-O-H 


i i 
H H 


Ai 


Meth}-] oxide 


Alcohol 



Polymerism. — Compounds having the same percentage com- 
position but different molecular weights are said to be poly- 



CAKBON DIOXIDE. 323 

meric. For example, N 2 4 is a polymer of N0 2 , and benzene, 
C 6 H 6 , is a polymer of acetylene, C 2 H 2 . 

Inorganic chemistry presents few examples of isomerism and 
polymerism: the best and most numerous examples are found 
among the compounds of carbon. Carbon appears to possess 
the property of forming complex and numerous compounds to 
a greater degree than other elements, and this property is 
in part explained by the theory that carbon atoms are linked 
to each other in such compounds. 

Carbon Dioxide, Carbonic Anhydride, C0 2 or 0=C=0, is 

a colorless gas with a slightly pungent odor and acid taste, and 
is commonly called carbonic acid. Carbon dioxide is formed 
when carbon or its compounds are burned in an excess of air 
or oxygen. Carbon in the form of charcoal burns with diffi- 
culty in oxygen freed from moisture by long contact with phos- 
phorus pentoxide, but it burns with great brilliancy in oxygen 
containing traces of moisture. Hence it appears that water 
plays an important part in the ordinary burning of carbon. 
The nature of the action of the water is not understood. 

In the formation of carbon dioxide 12 weights of carbon 
unite with 32 weights (2 volumes), of oxygen to form 44 
weights and 2 volumes of carbon dioxide. The gas density of 
carbon dioxide is 22. These are the data for the molecular 
formula C0 2 . 

Exp. 200. — a. Suspend a small piece of charcoal on the platinum wire 
over the platinum spoon of the apparatus, Fig. 75, p. lSo. Heat the char- 
coal intensely in a blast-lamp flame to expel volatile products, and 
allow to cool in a jar tilled with carbon dioxide. Fill the apparatus 
with oxygen, then drop in some phosphorus pentoxide and adjust the 
stopper holding the spoon. Allow the apparatus to stand a day or 
longer in order to free the gas as completely as possible from moisture 
by long contact with phosphorus pentoxide. Heat by a battery current 
the platinum wire which supports the charcoal. The charcoal will 
burn slowly as long as the wire glows, and will cease to burn as soon as 
the wire cools. 



324 THE FOURTH GROUP. 

b. Repeat the experiment with oxygen which has not been freed from 
water. The charcoal will burn brilliantly in the moist oxygen, and 
after the apparatus has cooled to the temperature of the room the 
volume of the gas will be the same as before the burning. 

Carbon dioxide is made on a large scale in the manufacture 
of hydrogen sodium carbonate by burning anthracite or coke. 
It is commonly generated for laboratory and other uses, when 
it is required nearly pure, by treating marble (calcium carbon- 
ate) with dilute hydrochloric acid: 

CaC0 3 + 2HC1 = CaCl 2 + H 2 + C0 2 . 

The gas thus prepared contains a trace of calcium chloride 
which is carried along with it mechanically, and also a little 
acid. By washing the gas with a solution of hydrogen sodium 
carbonate, and then passing it through a tube filled with cot- 
ton wool, these impurities may be almost completely removed. 
Pure carbon dioxide is best made by heating hydrogen sodium 
carbonate which decomposes as follows: 

2HNaCO, = C0 2 + H 2 + Na 9 CO B . 

Carbon dioxide is 1.524 times heavier than air, and hence 
may be collected by displacing air. The critical temperature 
of carbon dioxide is 30°. 9, below which point it condenses 
under pressure (at 0° under 36 atmospheres) to a colorless 
mobile liquid, boiling at —78°. When the liquid is exposed to 
the air a portion rapidly evaporates, and withdraws so much 
heat that the remainder is frozen to a white snow-like mass, 
which evaporates slowly. Solid carbon dioxide may be handled 
with impunity, as it is always surrounded by a gaseous layer 
which keeps it from actual contact with the skin. If, how- 
ever, it is pressed between the fingers, it blisters the skin like 
a hot iron. 

Carbon dioxide is one of the essential constituents of the 
atmosphere, which contains an average of 3 volumes of the 



CARBON DIOXIDE. 325 

dioxide in 10,000 volumes. It is supplied to the atmosphere 
by the respiration of animals, by fires, and by the decay of 
animal and vegetable matter. The leaves of plants absorb it 
from the air, and in the presence of sunlight decompose it, 
setting free oxygen, and using the carbon to form vegetable 
matter. Thus the plants prevent the increase of carbon dioxide 
in the atmosphere. 

Carbon dioxide has commonly been considered a poisonous 
gas, but its action is rather negative than positive. An in- 
crease in the amount of carbon dioxide in the air breathed is 
accompanied by a decrease in the amount of oxygen, thus pro- 
ducing oxygen starvation. Whenever the proportion of car- 
bon dioxide in the air breathed reaches 10 per cent, the oxygen 
being correspondingly diminished, asphyxia results ; and a 
diminution of the oxygen by one per cent may be followed by 
headache. 

Carbon dioxide sometimes accumulates in ]wells and coal 
pits, where it is known as choke-damp. Its presence^! may be 
detected by lowering a burning candle, which will go out if the 
air is unsafe to breathe. A candle will not burn in air con- 
taining 4 per cent of carbon dioxide, but such air may be 
breathed for a short time. 

Water dissolves its own volume of carbon dioxide gas at 15°, 
and at 0° 1.8 volumes. As the volume of a gas is inversely 
proportional to the pressure, so the quantity of the absorbed 
gas is proportional to the pressure. This is Dalton and 
Henry's law. According to this law, at double and triple the 
pressure water dissolves double and triple the amount of 
carbon dioxide. 

Carbonated water, now so extensively used as a beverage, is 
made by dissolving carbon dioxide in water under pressure. 
The solution of the gas in water is facilitated by the addition 
of a little common salt, owing to the reaction between carbonic 
acid and sodium chloride, whereby hydrochloric acid and 
hydrogen sodium carbonate are formed, thus: 



326 THE FOURTH GROUP. 

H„CO, 4- NaCl = HOI 4- HJSTaOO,. 






The effervescense of champagne, beer, soda-water, and other 
beverages is due to the escaping carbon dioxide. 

Tests for Carbon and Carbon Dioxide. — Carbon dioxide re- 
acts with a solution of calcium hydroxide to form an insoluble 
white precipitate of calcium carbonate : 

= C = + Ca<Q£ = = C<Q>Ca + H a O. 

This reaction is applied in testing for free carbon dioxide. 
If it is in combination as in carbonates, the latter are treated 
with acids to set the gas free. All carbon compounds yield on 
burning in air, oxygen, or with oxidizing substances carbon 
dioxide, which may be detected in the products of the com- 
bustion by means of a solution of calcium hydroxide. 

Exp. 201. — a. Burn charcoal in a jar of oxygen. Pour some lime 
water into the jar, and shake. The white precipitate which forms will 
dissolve with evolution of small bubbles of carbon dioxide on adding 
hydrochloric acid. 

b. Burn charcoal, a splinter of wood, a bit of paper, and some coal 

gas in a jar filled with air, and test the products of each combustion for 

carbon dioxide. 

Exp. 202. — Pour some clear lime water into a bottle fitted with 

tubes as in Fig. 87. Draw air into the lungs through the 

tube A. The air will bubble through the lime water, 

which will become slightly turbid owing to the carbon 

dioxide contained in the air. Next blow the air from 

the lungs slowly through B, and note the result. 

Exp. 203. — Place in the generator (Fig. 52, p. 54) lumps 
Fig. 87. Q f mar bi e (calcium carbonate), pour in some water, and 
then add hydrochloric acid until the gas escapes freely. Add more 
acid from time to time as required. 

Exp. 204. — a. Half fill a jar over water with carbon dioxide, then close 
the mouth of the jar with the hand and shake; open the jar under water. 
Repeat the shaking and opening under water as long as the gas is 
absorbed. Heat some of the carbonic acid water and note result, also 
test some of it by pouring it into lime water. 




CAKBONATES. 327 

b. Expose a moist blue litmus paper to carbon dioxide gas collected 
over water. It will turn red, showing that moist carbon dioxide reacts 
acid. Dry the paper; the carbon dioxide will escape, and the test paper 
will turn blue. Also test carbonic acid water with litmus. In case the 
blue color is not restored on drying, the water probably contains hydro- 
chloric acid carried over mechanically, and the experiment should be 
repeated with a fresh supply of water in the pneumatic trough. 

Exp. 205. — Connect the delivery-tube of a siphon bottle of carbonic 
acid water by means of rubber tubing with an inverted cylinder filled 
with and standing over water. On allowing the water to escape from 
the siphon bottle gas will collect in the cylinder, which can be tested 
for carbon dioxide. 

Exp. 206. — Into a jar filled with carbon dioxide thrust a burning 
splinter or taper. It will immediately be extinguished, as the gas does 
not support the combustion of carbon or compounds of carbon. 

Exp. 207. — Ignite in a lamp flame a piece of magnesium ribbon, held 
by forceps, and place it in a pint or more of carbon dioxide, The metal 
will burn brilliantly in the gas, with formation of white magnesium 
oxide and separation of carbon in form of a black powder. The mag- 
nesium oxide may be removed by dissolving in hydrochloric acid. 
Other metals, notably potassium and sodium, decompose carbon 
dioxide and set free carbon. 

Exp. 208.— a. Moisten the sides of a large glass jar with ammonia water, 
then pour in a little concentrated hydrochloric acid. A cloud of am- 
monium chloride will be formed. Pass into the jar a rapid stream of 
carbon dioxide. The cloud will indicate the upper surface of the gas. 
The gas in the jar may be poured into another vessel. 

b. Fill a jar in which there is no cloud with the gas, by displacing 
air ; then pour it into another jar, and test for the presence of the gas 
with a burning splinter, c. Dip a small jar into a large jar filled with 
the gas ; then take out the small jar, and test the gas in it as before. 



Carbonic Acid and Carbonates. 

The two possible hydroxyl derivatives of tetravalent carbon 

IKK ^OH , . *, 
110 Oil' a naeta-carbonic 



328 THE FOURTH GROUP. 

ATT 

acid, = C<qtt- These have not been isolated, but their 

salts and ethers are well known. A solution of carbonic an- 
hydride, C0 2 , reacts acid to test paper, and doubtless contains 
carbonic acid — whether the ortho- or meta-acid has not been 
determined. Few ortho-carbonates of metals are known, but 
the meta-carbonates are numerous and important. There are 
two classes of the latter : the acid meta-carbonates, as hydro- 

O TT 

gen potassium carbonate, = C<q_tt- and the normal meta- 

carbonates, as potassium carbonate, — C < q~jz- and calcium 
carbonate, = C<Q>Ca. 

O-Na 

Sodium Carbonate, Na 2 C0 3 or CO < Q _- N - , is manufactured 

on an enormous scale from common salt. Salt is converted 
into sodium sulphate by means of sulphuric acid. A mixture 
of sodium sulphate, coal, and calcium carbonate is heated in 
a flame oven. The coal reduces the sodium sulphate to 
sulphide, which at the same time reacts with the calcium 
carbonate to form sodium carbonate and calcium sulphide : 

Na 2 S0 4 -f 4C = JS> 3 S -f 4CO ; 
¥a s S -f CaC0 3 = &a 2 C0 3 + CaS. 

The fused mass of "black ash" thus obtained is treated 
with water, which takes up the sodium carbonate, leaving 
the calcium sulphide. The solution is evaporated, and the 
product is calcined in a flame oven to oxidize any sodium 
sulphide and to convert sodium hydroxide, which may have 
been formed by the action of lime, into carbonate. The 
crude sodium carbonate thus made constitutes the soda- 
ash of commerce. 

Crystallized sodium carbonate, Na 2 C0 3 -f- 10H 2 O, known 
as "sal-soda," is obtained by allowing a solution of soda-ash 



CARBONATES. 329 

to crystallize at a winter temperature, when clear rhombic 
prisms form. Crystals of the decahydrated salt melt in their 
water of crystallization at 34°, with formation of the mono- 
hydrate, Na 2 C0 3 + H 2 0, which is also formed as a white 
powder when sal-soda is exposed to air. 

Anhydrous sodium carbonate is obtained by drying crystals 
of sal- soda. It is a white powder which fuses at a red heat 
with slight loss of carbon dioxide. It is very soluble in water, 
and readily forms supersaturated solutions. Sodium carbonate 
reacts alkaline, and has the bitter taste of sodium hydroxide. 
It is extensively used in the manufacture of glass, sodium 
hydroxide, and hydrogen sodium carbonate. 

Sodium carbonate occurs as an efflorescence on the soil of 
desert regions. It is the alkali of the plains of the western 
United States. It occurs in some lake waters in sufficient 
abundance for profitable extraction. 

Exp. 209. — Dissolve 40 grams of soda-ash in 100 cc. of boiling water ; 
filter, and leave the solution in a glass jar. In case crystals do not 
appear after the "solution has stood a day orjonger, drop in a fragment 
of a crystal of sal-soda to start the growth of crystals. 

Exp. 210. — Determine the water of crystallization in sal-soda as 
follows. Weigh accurately a porcelain crucible ; then put into it a 
clean uneffloresced piece of sal-soda (about 2 grams), and weigh again. 
Set aside for a week, and then weigh. Finally, heat the crucible and 
contents moderately with a lamp until after repeated weighing the 
weight remains constant. The crucible should not be placed upon the 
pan of the balance until cool to the touch. Calculate the per cent of 
water lost at ordinary temperature, and also the total loss after heating. 
Compare the results obtained with the per cent of 9H 2 and 10H 2 O in 
Na 2 CO 3 + 10H 2 O. 

OH 

Hydrogen Sodium Carbonate, HNaC0 3 or C0< Q « , is com- 
monly known as bicarbonate of soda or baking soda. It 
separates as a white powder when a strong solution of sodium 
carbonate is exposed for a time to carbon dioxide, the bicar- 
bonate being less soluble than the carbonate. It is also formed 



330 



THE FOUKTH GEOUP. 



by the prolonged action of carbon dioxide on crystals of sal- 
soda : 



O-Na 



OH 



C0< 0lSa + 10H *° + 00 * = 2C0< 0-k + 9H 2 °' 



In the manufacture of bicarbonate of soda crystals of sal 
soda are placed in shallow layers in chambers through which 
carbon dioxide is passed for a number of days. The carbon 
dioxide is obtained from the decomposition of magnesite, a 
native magnesium carbonate, or from limestone by acids, and 
also by the burning of coke and anthracite coal. The water 
liberated by the reaction flows off saturated with sodium car- 
bonates, and the solution is evaporated to obtain the salt. 

Hydrogen sodium carbonate has a feebly alkaline reaction, 
and a less acrid taste than sodium carbonate. A solution of 
hydrogen sodium carbonate gives off carbon dioxide on boil- 
ing, and the dry salt is resolved by heating into normal car- 
bonate, carbon dioxide, and water. 






Few salts are so commonly used as bicarbonate of soda. 

Baking-powders are mixtures of 
it with hydrogen potassium tar- 
trate (cream of tartar), or some 
other salt. Baking-powder re- 
mains unchanged when dry, but 
on dissolving in water gives off 
carbon dioxide gas. When a 
9 mixture of flour and baking- 
powder and water is made, the 
gas set free inflates or raises the 
dough. One gram of hydrogen 
sodium carbonate yields, when 
decomposed by an acid, about 260 cc. of gas. Bicarbonate 




Fig. 88. 



CAKBOKATES. 331 

of soda is used in medicine as a mild anti-acid. With 
tartaric acid, and also with cream of tartar, it constitutes 
effervescing aperient powders. 

Exp. 211. — Place in the A, Fig. 88, some hydrogen sodium car- 
bonate ; heat, and allow the gas evolved to pass into lime water in the 
test-tube B. Note observations. 

Exp. 212. — Mix 15 grams of hydrogen potassium tartrate with the 
weight of hydrogen sodium carbonate required by the equation 

HKC 4 H 4 6 + HNaC0 3 = NaKC 4 H 4 6 -f C0 2 -f H 2 0. 
Hydrogen potassium Sodium potassium 

tartrate tartrate 

Observe whether any change occurs when the two salts are mixed, 
then place the mixture in a flask and pour in 100 cc. of water. Note 
the result. Heat to boiling, and filter the solution into a porcelain dish, 
and let stand for several days. Large crystals of sodium potassium tar- 
trate (Rochelle salt) will form. 

0— K 

Potassium Carbonate, K 2 C0 3 or CO < _ K , is the chief in- 
gredient of potash obtained by leaching wood ashes with 
water and evaporating the solution to dryness. The product 
is called potash, or potashes, because the lye from the ashes 
was formerly, and is in some localities at the present time, 
boiled down iu iron pots. The impurities in crude potash are 
potassium chloride, sulphate, and silicate, sodium salts, and 
other substances. In order to purify potash it is treated with 
a small quantity of hot water, which takes up the potassium 
carbonate and leaves most of the impurities undissolved. The 
solution is evaporated, and the residue is heated to redness. 
The purified product, from its pearly appearance, is known as 
pearl ash. Formerly wood ashes were the only source of 
potash, but in recent years large quantities have been obtained 
from beet-root molasses, wool washings, and potassium sul- 
phate. 

Pure potassium carbonate is best prepared by igniting hy- 



332 THE FOURTH GROUP. 

drogen potassium carbonate, a salt which can be easily ob- 
tained pure. 

Potassium carbonate is a white solid, which deliquesces in 
moist air to an oily liquid, and dissolves in less than its weight 
of water at ordinary temperature. It has a strong alkaline 
reaction and taste. It is used in the manufacture of other 
potassium salts, glass, and soft soap. 

Exp. 213. — Place half a pint to a pint of wood ashes upon a filter, and 
pour on an equal bulk of boiling water. After the water has mostly 
filtered through, pour it back again upon ihe ashes. Evaporate the 
solution obtained to dryness. Place a small portion of the residue in a 
test-tube, and add to it hydrochloric acid. An effervescence may be 
regarded as proof of the presence of a carbonate. To the remainder 
of the residue add a little water. A portion will remain undissolved. 
Test the reaction of the solution. 

OH 

Hydrogen Potassium Carbonate, HKC0 3 or CO < q £, is com- 
monly known as bicarbonate of potash, from the old formula, 
K 2 0.2C0 2 .H 2 0, carbonate of potash being in the same system 
K 2 O.C0 2 . It has also been called saleratus (aerated salt), 
since it is obtained by the action of an aeriform body upon a 
salt. It is prepared by passing carbon dioxide gas through a 
concentrated solution of potassium carbonate, when the less 
soluble hydrogen potassium carbonate will separate in crystals. 
It has a slightly alkaline reaction and saline taste. It dis- 
solves in about 4 parts of water at ordinary temperature. The 
solution loses carbon dioxide when boiled, and the dry salt is 
completely decomposed by heat into potassium carbonate, car- 
bon dioxide, and w T ater. Bicarbonate of potash was formerly 
used in cookery, but has been replaced by the cheaper bicar- 
bonate of soda. 

Commercial Ammonium Carbonate is manufactured by sub- 
liming a mixture of chalk (calcium carbonate) and ammonium 
sulphate or chloride, It is a compound of equal molecules of 



CARBONATES. 333 

hydrogen ammonium carbonate and ammonium carbamate, 
HNH 4 CO, + NH 2 NH 4 C0 2 . It is a colorless translucent fibrous 
mass, which on exposure to the air falls to a white powder of 
hydrogen ammonium carbonate, the carbamate slowly decom- 
posing and giving off ammonia. The commercial salt has a 
strong ammoniacal odor, and on this account is used in " smell- 
ing salts." It is used in medicine, and in the laboratory for 
precipitating carbonates. 

Ammonium Carbonate, (NH 4 ) 2 C0 3 + H 2 or co <oInH 4 + 

4 

H 2 0, is obtained as a crystalline powder when the commercial 
carbonate is digested for a time at 12° with aqueous am- 
monia. From a solution saturated at 30° to 35° it separates 
in transparent crystals. These lose ammonia, and change to 
the hydrogen ammonium carbonate. 

OH 

Hydrogen Ammonium Carbonate, HNH 4 C0 3 or C0<q_- n - h - , 

4 

is obtained when a concentrated solution of the commercial 
carbonate is saturated with carbon dioxide. It forms crystals 
which do not smell of ammonia nor lose their lustre in dry 
air. It is completely decomposed on warming into ammonia, 
carbon dioxide, and water. 

Calcium Carbonate, CaC0 3 or C0< Q >Ca, is the most 

abundant calcium compound in nature. It is dimorphous, 
occurring as calcite, crystallized in the hexagonal system, and 
having a density of 2.7, and as arragonite in the rhombic sys- 
tem, density 2.9. Limestone, marble, the shells of mollusca, 
egg-shells, coral, and chalk are chiefly calcium carbonate. It 
may be made by mixing solutions of an alkali carbonate and a 
calcium salt. If the precipitation is made at ordinary tem- 
perature the fine white powder which separates consists of the 
calcite form of crystals. But if the precipitation is made 
boiling hot the crystals are larger, and have the arragonite 
form. 



334 THE EOUKTH GBOTJP. 

Calcium carbonate is slightly soluble in pure water, 1 liter of 
cold or boiling water dissolving about 18 milligrams. It is 
more soluble in water containing carbonic acid. One liter of 
water saturated with carbonic acid dissolves at 10° 0.88 gram 
of calcium carbonate. Calcium carbonate separates on boil- 
ing, but the solution still contains 34 milligrams of it per 
liter. 

The carbonic acid solution of calcium carbonate loses carbon 
dioxide on exposure to air, and calcium carbonate precipitates. 
The natural waters of limestone regions contain calcium car- 
bonate, and the stalactites and stalagmites of caves are formed 
by the gradual deposition of it. The "hardness" of natural 
waters is chiefly due to calcium carbonate. 

Exp. 214. — To a dilute solution of calcium chloride add a solution of 
sodium carbonate so long as a precipitate forms. Heat to boiling, then 
collect the precipitate on a filter, and wash moderately. Test a portion 
of the precipitate for carbon dioxide. Represent by an equation the 
reaction between the calcium chloride and sodium carbonate. 

Exp. 215. — a. Dilute lime water with its bulk of water, and pass 
carbon dioxide through it until a clear solution results, b. Leave a 
portion of the solution in an open bottle, to find whether a precipitate 
forms after some days. c. Boil another portion until calcium carbonate 
separates, d. To some of the solution add a solution of soap, and ob- 
serve the formation of an insoluble lime soap. 

Strontium Carbonate, SrC0 3 or C0< o >Sr, occurs as stron- 

tianite, density 3.6 to 3.7. It is prepared from soluble stron- 
tium salts by the method described for the preparation of 
calcium carbonate. 

Barium Carbonate, BaC0 3 or C0< o >Ba, occurs as wither- 

ite, density 4.3. It is best prepared pure by precipitating a 
solution of barium chloride with ammonium carbonate. It is 
used in chemical analysis and for the preparation of barium 
salts. 



CAKBOHATES. 335 

The carbonates of the alkali-earth metals present a gra- 
dation of properties with the increasing atomic weights of the 
metals. Strontium carbonate is more soluble in water than 
calcium carbonate, and barium carbonate is more soluble than 
strontium carbonate. Calcium carbonate is converted into 
oxide at a red heat, and barium carbonate fuses and decom- 
poses only at very high temperatures. Strontium carbonate 
is intermediate in these properties between calcium carbonate 
and barium carbonate. 

Magnesium Carbonate, MgC0 3 or C0< Q >Mg, occurs as 

magnesite. This mineral does not lose carbon dioxide at 300°, 
and is but slightly acted upon by cold acids. Magnesium 
carbonate cannot be prepared by precipitating, a solution of a 
magnesium salt with an alkali carbonate, as the precipitate 
has the composition described under magnesia alba. It has 
been obtained with 3 and 5 molecules of water by allowing a 
solution of hydrogen magnesium carbonate to stand in a par- 
tially closed flask. 

Hydrogen Magnesium Carbonate, Mg(HC0 3 ) 2 or H0-C0-0- 
Mg-0-C0-0H, is not known in the solid state, but is probably 
formed when magnesia alba dissolves in water containing car- 
bon dioxide. Engel and Ville found that one liter of water 
containing carbon dioxide will dissolve at 

1 atmosphere and 19°. 5, . . . .25.79 grams of MgC0 3 . 
7.5 atmospheres and 19°. 5, . . . .51.2 " " 

7G5 mm. and 70°, 8.1 " " 

765 " " 100°, " " " 

The solutions contained very nearly one atom of magnesium 
to two molecules of carbon dioxide, the proportion required 
by the formula Mg(HC0 3 ) 2 . 

Magnesia Alba is a varying mixture of magnesium carbonate 



336 THE FOURTH GROUP. 

and hydroxide. It is made by adding sodium carbonate to a 
solution of a magnesium salt. The precipitate is washed with 
water and dried. It is also made by decomposing a solution 
of the hydrogen magnesium carbonate with a current of steam. 
Magnesia alba is almost insoluble in water, but dissolves readily 
in acids. It is used in medicine, and in the preparation of 
other magnesium compounds. 



Carbon Monoxide or Carbonic Oxide, CO, is formed when 
carbon dioxide is in contact with glowing carbon, as charcoal 
and coal : 

C0 3 + C = 2CO. 

2 vols. 4 vols. 

Hence it is produced -when carbon burns in an insufficient 
supply of air. 

Exp. 216. — a. Heat to redness in a gas combustion furnace or a coal 
fire fragments of charcoal contained in a piece of iron gas-pipe three 
quarters of an inch in diameter. The pipe may be conveniently bent 
in the form of a U tube if it is to be heated in a stove. Pass a slow 
stream of carbon dioxide into one end of the pipe. The gas which 
issues from the other end of the hot pipe will burn with the character- 
istic blue flame of carbonic oxide, b. Pass oxygen into the tube. Car- 
bonic oxide will be obtained as before. 

The experiment illustrates the changes which take place in 
an anthracite fire. The oxygen of the air entering at the 
grate bars combines with the carbon to form carbon dioxide, 
which passes upward through the glowing coals and is changed 
to carbonic oxide. If air is supplied over the coal the car- 
bonic oxide burns with a characteristic blue flame, providing 
the temperature is sufficient for its ignition. "When carbon 
burns only to the monoxide less than one third as much heat 
is obtained as when it burns to the dioxide ; thus, 12 grams 
of carbon in uniting with 16 grams of oxygen to form CO 



CARBONIC OXIDE. 337 

evolve 28,800 calories, while the same weight of carbon com- 
bining with 32 grams of oxygen to form C0 2 will evolve 96,800 
calories. 

Carbonic oxide may be obtained pure in a nnmber of ways. 
A good method is to heat oxalic acid and oil of vitriol together 
in a flask. The heat and oil of vitriol remove water from the 
oxalic acid, and liberate equal volumes of carbonic oxide and 
carbon dioxide : 

The carbon dioxide is then absorbed by a solution of sodium 
hydroxide or by milk of lime. 

Carbonic oxide is a colorless tasteless gas, with a peculiar 
feeble odor. It has a density of 14, which is little less than 
that of air. It is but slightly soluble in water. It condenses 
under a pressure of one atmosphere at —190° to a colorless 
transparent liquid which solidifies in vacuum at —211°. 

Carbonic oxide is very poisonous, small quantities in the air 
inhaled causing headache and insensibility. It unites with 
the coloring matter of the blood, forming a definite compound, 
which can be detected by means of the spectroscope. The 
characteristic spectrum of this compound reveals the cause of 
death in cases of poisoning with illuminating gas. 

Carbon monoxide burns in air or oxygen with formation of 
carbon dioxide. 

Exp. 217. — Fill a jar over water with carbon monoxide, and pour into 
the jar some lime water. If the gas is free from carbon dioxide the 
lime water will remain clear. Set fire to the gas, and when it has ceased 
burniug close the jar with the hand and shake. A white precipitate is 
evidence of the formation of carbon dioxide. 

It has been found that a mixture of perfectly dry carbonic 
oxide and oxygen is not ignited by a glowing platinum wire 
nor by the electric spark, and the following experiment shows 



338 THE FOURTH GROUP. 

that a flame of carbonic oxide is extinguished when thrust 
into dry air. 



Exp. 218. — The bottle D, Fig. 89, has a capacity of a litre or more. 
In order to dry the air in it, pour some concentrated sulphuric acid into 
the bottle, and allow it to remain corked some minutes. Pass carbonic 
oxide through a tube filled with fragments of potassium hydroxide, to 
free the gas from moisture, and then through the tube AB. Ignite the 
gas at B, remove the cork from B, and thrust the flame into D, placing 
the stopper C in the neck of the bottle. The flame will be extinguished 
if the air in the bottle is dry. Kepeat the experiment, using a bottle with- 
out attempting to dry the air. 

Traube explains the action of water in the burning of car- 
bonic oxide by the following equations : 

(1) CO + 2H 2 + 2 = CO(OH) 2 + H 2 2 . 

(2) H 2 2 + CO =CO(OH) 2 . 

(3) 2C0(0H) 2 = 2C0 2 +2H 2 0. 

Equation 1 represents the formation of carbonic acid and hy- 
drogen dioxide. In support of this view is the fact that 
hydrogen dioxide is obtained when carbonic oxide is burned 
from a jet on the surface of water. 

At a red heat carbonic oxide reduces many metallic oxides, 
and plays an important part in many metallurgical processes. 

Exp. 219. — Place some cupric oxide in a hard glass tube, and heat the 
part of the tube about the oxide. Pass into the tube carbonic oxide, 
and pass the gaseous product of the combustion into lime water. Note 
observations, and write the equation representing the reaction between 
the cupric oxide and carbonic oxide. 



CARBON" BISULPHIDE. 339 

CI 
Carbonyl Chloride, C0< cl . — This compound is formed 

when a mixture of equal volumes of dry chlorine and carbon 
monoxide is exposed to light. Carbonyl chloride at ordinary 
temperature is a suffocating gas. It is decomposed by water, 
with formation of carbon dioxide and hydrochloric acid : 

COCl 2 + H 2 = 2HC1 + C0 2 . 

The radical CO in combination is termed carbonyl. 

Carbon Bisulphide, CS 2 , is formed by the direct union of 
sulphur vapor with glowing charcoal. When charcoal is 
heated to the temperature of ignition in oxygen it continues 
to burn with evolution of sufficient heat to maintain it above 
the temperature of ignition. In the union of carbon and 
sulphur, on the contrary, heat is absorbed, and hence it is 
necessary to keep up the temperature by external heating. 12 
grams of carbon unite with 64 grams of sulphur with a ther- 
mal result of — 12,600 calories. 

Carbon disulphide is a colorless, mobile liquid, 'Which refracts 
light strongly, boils at 46°, and has an odor something like 
that of chloroform. Its vapor ignites at 149°, burning with 
a blue flame, with formation of carbon dioxide and sulphur 
dioxide. Large quantities of it are used to destroy vermin, as 
its vapor is poisonous. It is used in the arts as a solvent for 
caoutchouc, fats, and other substances. 

The commercial carbon disulphide has often a very disagree- 
able odor, due to impurities. It may be purified by distilling 
it from a flask filled with lumps of quick-lime. 

Thiocarbonic Acid, HoCS 3 . — Carbon disulphide reacts with a solution 
of sodium sulphide to form sodium thioearbouate, thus : 

CS, + Na,S = Na a CS s . 

Hydrochloric acid decomposes the thiocarbonate with separation of 
thiocarbonic acid, a yellow oil of disagreeable odor, which readily de- 



340 THE FOURTH GROUP. 

composes into hydrogen sulphide and carbon disulphide. A number 
of salts of the acid are known, but they are of little importance. It 
should be noticed that H 2 CS 3 is the analogue of H 2 C0 3 , which has not 
been isolated. 

Carbonyl Sulphide, COS, is formed when sulphur vapor and carbonic 
oxide are passed together through a hot tube. It is best obtained by 
other methods. It is a colorless gas, somewhat soluble in water, to 
whi" 7 r it imparts its peculiar odor and taste. It is supposed to exist in 
some sulphur waters. 



OH 

Carbamic Acid, C0<-«-tt .— When dry ammonia gas and 

carbon dioxide are mixed the two unite to form the ammonium 
salt of carbamic acid having the formula CO<^tt 4 . The 

j-N xl 2 

salt can be easily obtained in crystals by passing the dry gases 
into well-cooled absolute alcohol. Free carbamic acid is un- 
known. 

NH 

Carbamide or Urea, C0<jttt 2 - — This compound, like the 

preceding, is an amide * of carbonic acid. It is derived from 
the chloride of carbonic acid by the action of ammonia : 

CO<^} + 2NH 3 = CO<^ 2 + 2HC1. 

Urea exists to the amount of 2 or 3 per cent in human 
urine, from which it was originally obtained. In 1828Wohler 
discovered that it can be formed by heating an aqueous solution 
of ammonium cyanate, whose atoms rearrange themselves to 
form urea: 

= = N-NH 4 = 0=C<^ . 

This was an important discovery, as it was the first instance 
* For the general character of amides, see Compound Ammonias, p. 352. 



CYANOGEN COMPOUNDS. 341 

of the preparation of an organic compound from inorganic 
substances. 



Cyanogen Compounds. 

The univalent radical cyanogen, ON, exists in a large num- 
ber of compounds. Cyanogen is an acid-forming radical, 
somewhat similar in properties to the atom of chlorine. It 
was first obtained by Gay-Lussac in 1815, who, in his investi- 
gations of its compounds, was the first to show that a group of 
elements can act chemically as a simple element. It is one of 
the best examples known of a compound radical. Carbon and 
nitrogen do not unite directly, unless perhaps under the in- 
fluence of the induction spark. All nitrogenous organic com- 
pounds on ignition with metallic sodium yield sodium cyanide, 
and animal matter when heated with potassium carbonate 
yields potassium cyanide. Free nitrogen is absorbed by a 
glowing mixture of charcoal and potassium carbonate, with 
formation also of potassium cyanide. 

J\ ' I C N„ or i . — This compound is a color- 

Dicyanogen, ) 2 2 C=N 

less, highly poisonous gas, with a peculiar odor resembling that 
of peach kernels. It has a density of 26, corresponding to 
the formula C 2 N 2 . It condenses at low temperatures, or under 
pressure to a liquid which boils at — 20°. 7 and freezes to a 
crystalline mass melting at —34°. Cyanogen gas is best pre- 
pared by heating mercuric cyanide : 

ON CN 

Hg<™ = Hg + , n - 

The gas must be collected over mercury, as it is soluble in 
water and alcohol. It burns with a characteristic purple 



342 THE FOURTH GROUP. 

flame, with formation of carbon dioxide, nitrogen being set 
free. The aqueous solution of cyanogen soon undergoes 
change, with formation of hydrocyanic acid, oxalic acid, am- 
monia, carbon dioxide and urea, and deposition of a brown 
substance known as azulmic acid. Acids retard the decompo- 
sition. 

Exp. 220. — Heat dry mercuric cyanide in a hard glass tube and burn 
the escaping gas. The brownish-black residue remaining is para- 
cyanogen. 

Paracyanogen, (CU),. — This substance is obtained as de- 
scribed in the foregoing experiment. It has the same propor- 
tions of carbon and nitrogen as cyanogen, into which it is con- 
verted at 840°. Its molecular weight is unknown. 

Hydrocyanic Acid, HCN or H-C=N, is formed by the action 
of acids on cyanides. It is best obtained by pouring a cold 
mixture of 7 parts of oil of vitriol and 14 parts of water upon 
10 parts of coarsely powdered potassium ferrocyanide contained 
in a large retort. The neck of the retort is inclined upwards, 
so tnat only the more volatile portions will pass over during the 
distillation. In order to free the vapor from water it is passed 
through U tubes containing calcium chloride heated to 30°. 
The dry gas is condensed in a well-cooled receiver. In the 
preparation of an aqueous solution the vapors are cooled and 
received in water. 

Pure anhydrous hydrocyanic acid boils at 26°. 5. It is 
soluble in all proportions in water, alcohol, and ether. It has 
the odor of bitter almonds. Its aqueous solution is known as 
prussic acid. This decomposes on keeping into formic acid, 
ammonia, and other bodies. Traces of mineral acids retard 
this decomposition. Hydrocyanic acid is very poisonous, and 
very rapid in its action. The vapor when inhaled causes al- 
most instant death, and in small quantities produces headache 
and other troubles. An internal dose of -^ of a grain is 



CYANOGEN" COMPOUNDS. 343 

usually fatal to the human subject. The antidotes recom- 
mended are chlorine water and ammonia. 

Potassium Cyanide, KCN. — The pure salt is prepared by 
passing the vapor of hydrocyanic acid into an alcoholic solu- 
tion of potassium hydroxide when the salt separates in small 
crystals. Commercial potassium cyanide is made by melting 
dry potassium ferrocyanide. Potassium carbonate mixed with 
the ferrocyanide increases the yield, but the product is not as 
pure. 

Potassium cyanide is a white crystalline solid, very soluble 
in water. It is decomposed by acids, with evolution of hydro- 
cyanic acid. The salt smells of this acid, which is set free by 
the carbonic acid of the air. Hence potassium cyanide is 
kept in well-stoppered vessels. It is used in large quantities 
in electro-plating and in photography, and is a valuable re- 
agent in the laboratory. It is exceedingly poisonous. 

Silver Cyanide, AgCN, is similar in properties to the chlo- 
ride, bromide, and iodide of silver. It separates as a white 
precipitate when a solution of hydrocyanic acid is added to 
silver nitrate. It is insoluble in nitric acid, but easily soluble 
in ammonia, and is not changed by light. It dissolves readily 
in. potassium cyanide, and the solution on evaporation yields 
crystals of the double salt KAg(CN) 2 , which are soluble in 
four parts of water and are permanent in air. Mention has 
already been made of the use of this compound in electro- 
plating. 

Cyanides of Copper. — Potassium cyanide produces in a cold 
solution of copper sulphate a yellowish precipitate of cupric 
cyanide, Cu(ON),, which soon loses cyanogen and changes 
into Cu 2 (CN) 3 .Cu(CN") 2 . This on heating changes to 
cuprous cyanide, which is white. The cyanides of copper are 
insoluble in water, but dissolve readily in potassium cyanide. 
Such solutions are used in electro-plating with copper. 



344 THE FOURTH GROUP. 

Trihydrocyanic Acid, H 3 C 3 N 3 . — This compound is formed 
when an aqueous solution of hydrocyanic acid, to which a few 
drops of potassium hydroxide have been added, is allowed to 
stand. It fuses at 180°, and at higher temperatures is re- 
solved into hydrocyanic acid gas. 

Cyanogen Chloride, CNC1, and Cyanuric Chloride, C 3 N 3 C1„ 

possess gas densities corresponding to their formulas. 

Normal Cyanic Acid, N-C-OH, is unknown in the free state, 
but there are ethers of it, as for example methyl cyanate, 
N=C-0-CH 3 . 

Isocyanic Acid, = C=N-H, is also called carbimide.* It 
is best prepared by heating cyanuric acid: 

CA(OH), = 30CNH. 

The vapors are passed into a vessel surrounded by a freezing 
mixture. On removing the liquid obtained from the freezing 
mixture it becomes hot, and changes to a porcelain-like mass 
of cyamelide, a compound of unknown molecular weight. 

Potassium Isocyanate, = C = N-K, is formed when potas- 
sium cyanide is melted in air, or better with some oxide such 
as manganese dioxide or red lead. Cyanates do not yield 
cyanic acid when treated with an acid, but are decomposed 
with evolution of carbon dioxide and formation of an ammo- 
nium salt, thus: 



= C=N-K 4- 2HC1 4- H„0 = KC1 + NH.C1 + CO 



Exp. 221. — Fuse in a glass tube or an iron spoon a little potassium 
cyanide to which some red lead has been added. When the mass has 

* An imide is a compound derived from ammonia by the replacement 
of two atoms of hydrogen by a bivalent acid radical, and hence contains 
the group NH, 



CYANOGEN" COMPOUNDS. 345 

cooled dissolve some of it in water, and under a hood with a good 
draught add hydrochloric acid to acid reaction, and heat the solution 
to boiling. Add an excess of potassium hydroxide to the acid solution 
and boil again; ammonia will escape freely. Impure potassium cya- 
nide containing a little cyanate with similar treatment gives a little 
ammonia, but not as much as obtained from the cyanate. 

Cyanuric Acid, C 3 N 3 (0H) 3 , crystallizes from an aqueous 
solution with two molecules of water. It results from the 
action of water on cyanuric chloride: 

C 3 N 3 C1 3 + 3HOH = 0,N,(OH), + 3H01. 

It is best obtained by passing chlorine over melted urea 
when the following reaction takes place: 

3CO(NH a ) a + 3Cl = N + H01-f2NH 4 01 + 3 ¥ 3 (OH) 3 . 

Thiocyanic or Sulphocyanic Acid, NeeC-S-H, is a colorless 
liquid with an odor similar to that of strong acetic acid. It 
is obtained, anhydrous when dry hydrogen sulphide is passed 
over the mercuric salt. Thiocyanates are formed by the direct 
union of sulphur with a cyanide, as when potassium cyanide 
and sulphur are fused together. The ammonium salt is 
formed when prussic acid is warmed with yellow ammonium 
sulphide: 

(NH 4 ) a S„ + HON = NH 4 SH + NCSNH 4 . 

Ammonium thiocyanate is best prepared as follows. A mix- 
ture of 8 parts of carbon disulphide, 30 parts of alcohol, and 
30 parts of concentrated ammonia is allowed to stand until 
the carbon disulphide is dissolved, when the solution is con- 
centrated. It yields on cooling crystals of the salt. The 
changes which occur in this process are complicated. The 
ammonia and carbon disulphide react to form intermediate 
products, which are decomposed, when the solution is heated, 
with formation of ammonium thiocyanate. 



346 THE FOURTH GROUP. 

Ammonium or potassium thiocyanate is used in testing for 
ferric salts, to whose solutions it imparts a red color, ferrous 
salts giving no coloration with the reagent. 



Methane and Derivatives. 

Methane or Marsh Gas, CH 4 . Gas density, 8. — Methane is 
the simplest compound of carbon and hydrogen known, and 
is, next to hydrogen, the lightest known gas. It is formed in 
marshes by the slow decomposition of vegetable matter under 
water. It is the "fire-damp" of coal mines, and is contained 
in the mixture of gases from natural-gas wells. It may be 
obtained synthetically by passing a mixture of carbon disul- 
phicle and hydrogen sulphide over glowing copper: 

CS 2 + 2H 2 S + 8Cu = CH 4 + 4Cu a S. 

Hence methane may be formed from substances which are 
not products of animal or plant growth. Methane is com- 
monly prepared by heating a mixture of sodium acetate and 
sodium hydroxide: 

CH 3 COOXa + ]S T aOH = CH 4 + tfa a CO,. 

Exp. 222. — Heat an intimate mixture of 1 part of sodium acetate and 
4 parts of soda-lime (a mixture of sodium hydroxide and lime) in a hard 
glass tube or other suitable vessel. Collect the gas over water. Lift a 
jar of the gas from the water, and before turning the mouth of the jar 
upwards ignite the gas. Mix 1 volume of the gas with 10 volumes of 
air and ignite. Mix 1 volume of the gas with 2 volumes of oxygen in 
a stout tube of 50 cc. capacity. Wrap the tube in a towel for safety, and 
ignite the gas. 

Methane made from sodium acetate contains a few per cent of hydro- 
gen and a little ethylene, C 2 H 4 . The latter may be removed by passing 
ihe gas through a tube filled with pumice stone drenched with strong 



METHANE. 34? 

sulphuric acid. The presence of ethylene increases the luminosity of 
burning methane. Methane from sodium acetate after remaining in a 
gas-holder several days is sufficiently pure for experiments. The water 
in the gas-holder slowly absorbs the impurities, which add to the light of 
the burning gas. 

Methane is colorless and odorless, and but slightly soluble 
in water. Mixed with sufficient air or oxygen, it burns with 
explosive violence, with formation of carbon dioxide and water. 

The four hydrogen atoms of methane may be successively 
replaced by chlorine aided by sunlight, with formation of 
chlorine substitution products of methane, thus: 

CH 4 + C1 2 = OH3CI + HCI; 

CH3CI + C1 2 = CH 2 C1 2 + HC1; 

0H 2 C1 2 + 01, = CHC1 3 + HOI; 

0HC1 3 + Cl 2 = 0C1 4 + HOI. 

The reactions are of theoretical interest, but in practice the 
substances are obtained by other methods. 

The group remaining after the replacement of part of the 
atoms of a molecule is frequently called a "residue" or a 
"rest:" the latter is the German term. Such a group is ob- 
viously a compound radical. For example, by substituting an 
atom of chlorine for one of hydrogen in CH 4 we have CH 3 01, 
and CH 3 is a residue or radical. It is called methyl, and 
0H.C1 methyl chloride. 

To indicate the derivation of methyl chloride from methane 
it is also called monochlormethane. The other chlorine sub- 
stitution products for the same reason are termed di-, tri-, and 
tetra-chlormethane. 

Methyl Chloride, CH S C1, is now obtained in large quantities 
as a by-product in the manufacture of beet sugar. It boils at 
—23°. The pressure of its vapor at ordinary temperature is only 
4 or 5 atmospheres. Methyl chloride is used for producing 



348 THE FOURTH GROUP. 

low temperatures by evaporation, and in the manufacture of 
aniline dyes. 

Methyl Iodide, CH 3 I, is a colorless liquid boiling at 44°. It 
has an ethereal odor, and partially decomposes in light with 
separation of iodine. It is prepj^ed by adding iodine gradu- 
ally to methyl alcohol and phosphorus: 

10CH.OH + 101 + 2P = 10CH 3 I+2PO(OH) 3 + 2H 2 O. 

Chloroform or Trichlormethane, CHC1 3 , is obtained by dis- 
tilling a mixture of chloride of lime, common alcohol, and 
water. It has an agreeable odor, and a sweetish, burning 
taste. It boils at 61°, is almost insoluble in water, but mixes 
in all proportions with alcohol, ether, and other organic liquids. 
It is a solvent for bromine, iodine, phosphorus, and many 
organic compounds. The vapor of chloroform when inhaled 
produces insensibility; hence its use in surgery. 

Methyl Alcohol, CH 3 0H, is one of the products of the de- 
structive distillation of wood. The wood alcohol or wood spirit 
of commerce is a mixture of methyl alcohol, acetone (0 3 E 6 0), 
water, and other compounds. It is used in place of common 
alcohol as a solvent and for burning, and is used in the prep- 
aration of methyl compounds. Its vapor when inhaled for 
some time causes headache. In England it is mixed with 
common alcohol to render the latter unfit for use in the manu- 
facture of beverages. The object of this is to avoid the tax 
which is placed on common alcohol. 

Pure methyl alcohol is obtained by the decomposition of 
certain methyl compounds. It boils at 55°.l, mixes in all 
proportions with water, and burns with a pale-blue flame. It 
is similar to common alcohol in chemical properties. 

Formaldehyde, HCHO, is the first of a series of compounds 
known as aldehydes which are formed from alcohols by the 



CONSTITUTION" OF DERIVATIVES OF METHANE. 349 

removal of two atoms of hydrogen by partial oxidation. 
Methyl alcohol is converted into aldehyde by the following 
reaction: 

CH 3 OH + = HCHO + H a O. 

Formaldehyde is an unstable compound, uniting with 
oxygen to form formic acid. 

Formic Acid, HCOOH, is so called because it was first ob- 
tained by distilling red ants {Formica mifra). It can be pre- 
pared in a number of ways — best, however, by heating oxalic 
acid with glycerine, when a mixture of formic acid and water 
distils over. The following reactions by which formates are 
produced synthetically are interesting: 

(1) CO + HOK = HCOOK. 

(2) 2C0gg + 2K = CO°g + CO° K + H 2 0. 

The first reaction takes place best at 200°, and the second 
occurs when potassium is suspended in moist carbon dioxide. 

Anhydrous formic acid boils at 99°. 9, forms crystals on 
cooling which melt at 8°. 6. It has a sharp acid odor, and is 
painfully corrosive to the skin. It decomposes completely on 
warming with sulphuric acid into carbon monoxide and water. 



Constitution of the Derivatives of Methane. 

The first question regarding the constitution of methane is: 
Do the four atoms of hydrogen differ from each other in their 
relation to the atom of carbon ? No satisfactory answer can 
be given: we can only say that there is as yet no evidence of 
any difference in the hydrogen atoms in OH 4 . Methyl chlo- 



350 THE FOURTH GROUP. 

ride, CH 3 C1, has been prepared in a number of different ways, 

and the different preparations have been found to possess 
identical properties. If there were a difference in the hydrogen 
atoms in CH 4 we might expect to obtain two or more isomeric 
methyl chlorides. We therefore consider all of the four 
hydrogen atoms -to be similarly related or joined to the atom 
of carbon, and represent this view by the graphic formula 

H 

H-C-H. 

A 

If one atom of hydrogen in methane is replaced by another 
univalent element, it can make no difference which one is 
replaced, since but one arrangement or linking of the atoms 
is possible. For example, methyl iodide may be represented 

by 

H H 

H-C-I; H-C-H; H.CI; or CH.I. 

i i 

H I 

All the formulas indicate that three atoms of hydrogen and 
one of iodine are joined to one atom of carbon. 

Methyl iodide is obtained from methyl alcohol, and, as 
shown later, the alcohol can be obtained from the iodide. 
Since methyl iodide contains the radical OH 3 , we infer that 
methyl alcohol also contains the same radical, and that in 
methyl alcohol three atoms of hydrogen are linked to carbon. 
What relation does the atom of oxygen and the fourth atom 
of hydrogen bear to the carbon? The reactions by which 
methyl alcohol is formed from compounds of known consti- 
tution, and the reactions of the alcohol with other compounds, 
answer the question. Methyl iodide and potassium hydroxide 
react to form methyl alcohol: 



CONSTITUTION OF DEEIVATIYES OF METHANE. 351 

H3C-I+K-OH = H3C-OH + K-I. 

By acting on methyl alcohol with hydrogen chloride, OH 
can be replaced by 01 by a reaction analogous to that which 
occurs between potassium hydroxide and hydrogen chloride, 
thus: 

H 3 C-0H + H-Cl = H 3 C-C1 + H 2 0; 

K-0H + H-C1 = K-C1 + H 2 0. 

We conclude, then, that methyl alcohol contains the radicals 

H 

CH 3 and OH, and that its graphic formula is H-C-O-H. 

i 
H 

Since formaldehyde is derived from methyl alcohol by the 
removal of two atoms of hydrogen, and is converted into formic 
acid by the addition of one atom of oxygen, the natural infer- 
ence is that the aldehyde contains the hydroxyl group which 
exists in both the alcohol and acid. This view is represented 
by the following formulas: 



H 


H 


H 


H-C-OH; 


-C-OH; 


0=C-OH. 


i 
H 


i 




Methyl alcohol 


Formaldehyde 


Formic acid 



Aldehydes, however, do not exhibit the deportment of 
hydroxides. When acetaldehycle, for example, is treated with 
a chloride of phosphorus, the compound CH 3 CC1 is not formed, 
as would be the case if hydroxyl was replaced by chlorine: but 
the atom of oxygen is replaced by two atoms of chlorine, with 
formation of the compound CH 3 CC1 2 H, known as diehlore- 
thane. If the aldehydes do not contain hydroxyl, we cannot 

II 
represent formaldehyde by A _qtt The other possible for- 

i 



352 THE FOURTH GROUP. 

TT 

mula is I If the constitutional formula of formalde- 

hyde is written H-CHO, no view is expressed regarding the 
structure of the univalent group -CHO. This group is char- 
acteristic of the class of compounds known as aldehydes, which 
have the general formula R-CHO, R representing a univalent 
hydrocarbon radical. Thus, replacing R by methyl, 0H 3 , we 
have acetaldehyde, CH 3 -CHO. 

Aldehydes unite directly with oxygen to form acids: 

HCHO + = HCOOH; 
CH 3 CHO + = CH3COOH. 

Formic acid has the empirical formula CH 2 2 . Since it is 
monobasic, containing but one atom of hydrogen replaceable 
by basic radicals with formation of salts, the two atoms of 
hydrogen do not bear the same relation to the atom of carbon. 
One hydrogen atom is supposed to be linked to carbon by one 
atom of oxygen. In other words, formic acid, in common 
with oxygen acids in general, contains hydroxy]. Since the 
other atoms of oxygen and of hydrogen are differently related 
to the carbon atom, each must be linked directly to the car- 

H-C-O-H 

bon atom, giving the structural formula || for formic 

& O 

acid. The univalent group -COOH, containing the radicals 

carbonyl and hydroxyl, is called carboxyl. Later we shall 

learn that there are good reasons for supposing acetic acid to 

contain the carboxyl group, and to have a constitution similar 

to formic acid. 



Compound Ammonias. 

The hydrogen of ammonia may be partly or entirely re- 
placed by hydrocarbon radicals with formation of a class of 



COMPOUND AMMONIAS. 353 

bodies known as amines. If the replacing radical contains 
oxygen, and is acid in character, the resulting compound 

ATTT 

ammonia is called an amide. Carbamide, CO<-^-rr% ma y De 

viewed as formed by the replacement of hydrogen in two 
molecules of ammonia by the acid radical carbonyl, GO. 
Acetamide, CH 3 CONH 2 , contains the acid radical acetyl, 
CH3CO, described later. 

When methyl iodide is heated with an alcoholic solution of 
ammonia, the following reactions occur : 

NH 3 + CH3I = NH 2 CH 3 + HI. 

Methylamine 

The methylamine exchanges an atom of hydrogen for CH 3 : 
NH 2 CH 3 + CH 3 I = NH(CH 3 ) 2 + HI. 

Dimethylamine 

The reaction continues until all of the hydrogen is replaced 
by methyl: 

NH(CH,) a + CH,I = N(CH 3 ) + HI. 

Trimethylamine 

The trimethylamine combines with methyl iodide to form 
tetramethylammonium iodide: 

N(CH 3 ) + CH 8 I = N(OH,) 4 L 

The hydriodic acid of the reactions combines with the ammonia 
used and the amines to form iodides. 

Methylamine, NH 2 CH 3 , is a gas which liquefies at a few 
degrees above zero. It closely resembles ammonia in chemical 
properties and odor. It is more soluble in water than ammonia. 
Like the latter, it forms salts by direct union with acids. 



354 THE FOUBTH GKOTTP. 

For example, with hydrochloric acid it forms methylammo- 
nium chloride, NH,CH 8 C1. 

Dimethylamine, NH(CH 3 ) 2 , is very similar to the preceding 
compound. It boils between 8° and 9°. 

Trimethylamine, N(CH 3 ) 3 , is not uncommon in nature, 
being found in a number of plants, in animal liquids, and 
especially in herring brine. It is obtained in quantity in the 
distillation of the " vinasses" or the waste liquids of the beet- 
sugar refineries. It boils between 9° and 10°, and has a 
strong, fish-like odor. It is used in the arts and in medicine. 

Tetramethylammonium Iodide, N(CH 3 ) 4 I, is the chief prod- 
uct of the reaction between methyl iodide and ammonia. It is 
a bitter- tasting, white crystalline salt. When freshly prepared 
silver oxide is added to its aqueous solution tetramethyl- 
ammonium hydroxide, N(CH 3 ) 4 OH, is formed. This last com- 
pound is a crystalline solid, which is similar in properties to 
the fixed caustic alkalies. It forms salts with acids which are 
not decomposed by potassium hydroxide. 

It will be observed that tetramethylammonium hydroxide 
is ammonium hydroxide with four atoms of hydrogen replaced 
by four methyl groups. One of the reasons for the formula 
NH 4 0H is the existence of compounds like N(CH 3 ) 4 OH. 



Derivatives of Ethane. 

These contain two atoms of carbon linked directly to each 
other, and are often called dicarbon compounds in distinction 
from derivatives of methane or monocarbon compounds. There 
are a number of ways by which dicarbon compounds can be 
formed synthetically from molecules containing only one atom 
of carbon. Such reactions are of the highest importance, 



DEEIVATIVES OF ETHANE. 355 

since they are the means not only of preparing complex com- 
pounds, but also give knowledge of their constitution. 

Ethane, C 2 H 6 or H 3 C-CH 3 . Gas density, 15.— This gas 
accompanies petroleum, and occurs in the gases of the natural 
gas wells. It is formed when methyl iodide and zinc are 
heated together at 150° : 

2CH 3 I + Zn = C 2 H 6 + Znl 2 . 

The zinc removes the iodine from two molecules of methyl 
iodide, and two methyl radicals unite to form one molecule of 
ethane. This synthesis leads to the constitutional formula 

HH 
H-C-C-H. 

ii 

Ethane has also been obtained in other ways, and the vari- 
ous preparations have not been found to differ in properties. 
Hence the inference is that there are no isomers of C 2 H 6 . A 
different linking of the atoms than the one given is not con- 
ceivable in the present state of chemical science. 

Chlorine replaces successively all of the hydrogen in ethane. 
The first stage of the reaction may be represented as follows: 

HH HH 

H-C-C-H + C1 2 = H-C-c'-Cl + HOI, 
H H II II 

CH 3 CH 2 C1 is ethyl chloride. It is prepared by other 
methods, one of which is by acting on ethyl alcohol with 
hydrogen chloride: 

C a H-OH+II01 = C a H s -Cl + H 8 0. 



356 THE FOURTH GROUP. 

Ethyl chloride boils at 12°. 5. Ethyl bromide is similar to the 
chloride, and boils at 38°. 4. 

Methyl and ethyl belong to a class of hydrocarbon radicals 
which exist in alcohols, and snch radicals are sometimes 
termed alcohol radicals. Thus, for example, in propyl alco- 
hol, C 3 H 7 OH, the alcohol radical is C 3 H 7 . , 

Ethyl Iodide, C 2 H 5 I. Gas density, 78. — This compound is 
obtained by acting on ethyl alcohol with iodine and phos- 
phorus, the reaction being analogous to that given for the for- 
mation of methyl iodide. Ethyl iodide is a colorless liquid, 
boiling at 72°, and having a density of 1.93. It is soluble in 
alcohol and ether, but almost insoluble in water. Like methyl 
iodide and many other organic iodides, it decomposes on ex- 
posure to light, with separation of iodine, which colors the 
liquid red. 

It is especially adapted for the synthesis of ethyl com- 
pounds, as it readily exchanges iodine for metals and com- 
pound radicals. Experience has shown that iodine in organic 
compounds is more readily replaced than chlorine or bromine. 

Exp. 223.— Place in a flask of about 800 cc. capacity 200 cc. of ordi- 
nary alcohol and 20 grams of red phosphorus, and then add gradually 
200 grams of iodine. The mixture will become warm, and it may be 
necessary to cool it by allowing water to flow over the flask. After all 
of the iodine has been added, connect the flask with an upright con- 
denser and allow to stand for a day. Then distil off the ethyl iodide, 
best by partly immersing the flask in a water-bath. The distillate con- 
tains ethyl iodide, water, alcohol, and a little hydriodic acid. The im- 
pure iodide is washed with water containing a little sodium hydroxide 
to remove the alcohol and acid, the former being taken up by the water 
and the latter by the alkali. In order to dry the iodide it is left in 
contact with calcium chloride for a day, without exposure to light. 
Finally, the ethyl iodide is poured off from the calcium chloride and 
distilled. For many purposes it is not necessary to dry and redistil 
the iodide before using. 

Leave a small portion of the ethyl iodide exposed to light, and keep 
the remainder out of light. 



ALCOHOL. 357 

Ethyl Alcohol, C 2 H 5 0H. — In the same way that methyl alcohol 
is derived from methyl iodide, ethyl alcohol may be derived 
from ethyl iodide. In place of potassium hydroxide 15 parts 
of water at 100° will answer : 

C 2 H 5 I + HOH = C 2 H 5 OH + HI. 

The radical ethyl has been shown to have the structure 
H H 
H-C- C-. The above reaction leads then to the structural 

i h 

H H 
formula H-C-O-OH for ethyl alcohol, in which one hydrogen 

H H 

atom is linked by oxygen to one carbon atom; that is, the 
alcohol contains the hydroxyl group, and is ethyl hydroxide. 
Proof of the existence of the radical ethyl is found in a large 
number of the reactions of the alcohol by which ethyl is trans- 
ferred to other combinations. Admitting the existence of 
ethyl in common alcohol, then the above structural formula 
is the only one possible. It is more conveniently written 
CII 3 CH 2 OH. It will be noticed that ethyl alcohol differs from 
methyl alcohol by CH 2 . There is the same difference between 
ethane, 2 H 6 , and methane, CH 4 , and many of the derivatives 
of these hydrocarbons. 

Ethyl alcohol is the best known of the alcohols, and is com- 
monly called alcohol, and also, more rarely than formerly, 
spirits of wine. It is the basis of intoxicating beverages, which 
have been manufactured in all ages and by all nations from 
cereals and fruits. The juices of fruit contain sugar, which is 
converted into alcohol by fermentation. The cereals and pota- 
toes are largely composed of starch, which is first transformed 
into a sugar, and this in turn yields alcohol when fermented. 
The alcohol is separated from the non-volatile products of the 
fermented liquid by distillation. The crude spirit obtained 



358 THE FOURTH GKOUP. 

from cereals or potatoes contains a large proportion of water 
and small quantities of a number of organic compounds. By 
repeated rectification and filtering through charcoal a spirit is 
obtained containing 90 to 96 per cent of pure alcohol, the re- 
mainder being water. All the water cannot be separated from 
alcohol by distillation. To remove the water in common 
alcohol, the latter is left in contact for some time with lime, 
potassium carbonate, or some other compound Avhich takes up 
water, absolute alcohol being finally distilled off. 

Absolute alcohol boils at 7S°.4, and has a density at 15° of 
0.794. It is very hygroscopic, and mixes in all proportions 
with water. It burns with a pale-blue flame. Next to water, 
alcohol is the most common solvent, dissolving fats, oils, 
resins, and other kinds of carbon compounds, and also some 
inorganic salts. 

Alcohol is used for a great variety of purposes in the in- 
dustrial arts. It is indispensable in the preparation of many 
medicines, and in the laboratory. 

Aldehyde or Acetaldehyde, CH 3 CH0, is derived from alcohol 
by taking out two atoms of hydrogen: 

CH 3 CH 2 OH + = CH 3 CHO + H 2 0. 

Various oxidizing agents effect the change. Aldehyde is 
also formed when a mixture of sodium formate and sodium 
acetate is heated, thus: 

OH s OOONa = CH s CHO + N aj CO s . 

Both reactions serve for the preparation of other aldehydes. 

Acetaldehyde boils at 20°. 8, giving off a vapor with a 
peculiar irritating odor. It mixes in all proportions with 
water and alcohol. Small quantities of inorganic acids convert 
aldehyde into a polymeric compound called paraldehyde, 



ACETIC ACID. 359 

C 6 H I2 3 , boiling at 124°, and possessing a gas density corre- 
sponding to the molecular formula given. Aldehyde readily 
unites with oxygen to form acetic acid. A mixture of alde- 
hyde, ammonia, and silver nitrate deposits a mirror of silver 
on warming. The reaction serves to detect traces of aldehyde. 

Acetic Acid, CH 3 C00H. — This di-carbon compound can be 
formed synthetically from a monocarbon compound as follows. 
Methyl iodide and potassium cyanide react when heated to- 
gether to form potassium iodide and methyl cyanide or 
acetonitrile, thus: 

H 3 CI + K-C=N = H,C-C=N. 

Acetonitrile on warming with potassium hydroxide yields 
acetic acid and ammonia, thus: 

H 3 C-C=N + 2H 2 = H 3 0-C0 2 H + NH 3 . 

It will be observed that the constitution thus far derived for 
acetic acid is based upon that assigned to methyl cyanide, in 
which it is assumed that the two carbon atoms are directly 
linked. This view is supported by the deportment of methyl 
isocyanide,* H 3 C-N"=C, which is decomposed with difficulty by 
alkalies, but is converted by acids into two monocarbon com- 
pounds, thus: 

H 3 C-N=C + 2H 2 = H 3 0-NH 2 + HCOOH. 

The difference in deportment of the two methyl cyanides is 
best explained by the theory that the carbon atoms in one are 
linked together and are not in the other. If the structural 
formula of acetonitrile is correct, then in all probability the 
two carbon atoms are linked together in acetic acid, and we 

* Methyl isocyanide is obtained by acting on methyl iodide with 
silver cyanide. It boils at 58' to 59 . Methyl cyanide boils at 82°. 



360 THE FOURTH GROUP. 

must next discuss the relation of the four atoms of hydrogen 
and two of oxygen to the carbon atoms. The radical CH 3 
was transferred from CH 3 I to CH 3 CN, and we may assume 
that CH 3 remains unchanged when methyl cyanide is con- 
verted into acetic acid. The fourth atom of hydrogen and 
the two atoms of oxygen and one atom of' carbon form a group 
united to CH 3 . Regarding carbon as tetravalent and oxygen 
as bivalent, we see, if our theory thus far is correct, that 
there are two possible formulas for acetic acid, viz. : 

H HO 

(1) H-6-o/i- (2) H-C-C-O-H. 
H H x H 

It has already been shown that oxygen acids when treated 
with phosphorus trichloride or pentachioride exchange hy- 
droxy! for chlorine. Acetic acid by similar treatment yields 
acetyl chloride, thus: 

3CH 3 C0 2 H + PC1 3 = 3CH 3 C0C1 + P(OH) 3 . 

And acetyl chloride with water yields acetic acid: 



CH 3 CO 1 CI + H 1 OH = CH 3 COOH + HCL 

These reactions show that acetic acid contains one hydroxyl 

group, and that its properties are best represented by the 

formula 

HO 
i ii 
H-C-C-O-H. 



The shorter formula CH 3 COOH embodies the same view. 
Acetic acid, as its formula indicates, is monobasic. Acetyl, 
0H 3 CO, is the acid radical of acetic acid. 



ACETIC ACID. 361 

Alcohol is converted by oxidizing agents first into aldehyde 
and then into acetic acid. Alcohol is unchanged by pure air un- 
less in contact with platinum black; but fermented liquors, 
containing nitrogenous matter, become sour on exposure to air 
unless too large a proportion of alcohol is present. The sour 
product is vinegar, which is a very dilute solution of acetic 
acid, and contains small quantities of other compounds, im- 
parting color and flavor. The formation of vinegar from 
dilute alcoholic solutions is due to the growth of a microscopic 
organism (Mycodermi aceti), which in some way serves as a 
carrier of atmospheric oxygen. Vinegar is manufactured from 
many different substances, such as poor wines, cider, and 
malt. The formation of vinegar is hastened by addition of 
"mother of vinegar," which contains Mycodermi aceti, and 
also by allowing the liquor to flow through a mass of shavings 
in which air circulates. 

Another considerable source of acetic acid is the liquid prod- 
uct of the distillation of wood. This contains, among other 
compounds, acetic acid. The acid is neutralized with lime or 
soda, and the salt, which remains after evaporation, is heated 
to remove impurities. The sodium or calcium acetate is 
mixed with sulphuric acid, and the acetic acid is distilled off. 

Pure acetic acid is a colorless liquid, boiling at 119°. Its 
gas density is 45 at 125°. The density gradually diminishes 
with increasing temperature, and at 250° is 30, the density 
corresponding to the formula C 2 H 4 2 . Acetic acid when 
cooled forms large transparent crystals melting at 16°. 7. The 
presence of a small quantity of water lowers the melting point 
considerably. The acid which crystallizes on cooling is known 
as glacial acetic acid. It cannot be separated from dilute 
solutions by distillation, but is prepared by acting on anhy- 
drous sodium acetate with sulphuric acid: 

2CH 3 COONa+S0 2 <^ = 2CH 3 COOH + SO,<^X 



362 THE FOURTH GROUP. 

The acetic acid which is distilled off contains a little water. 
To prepare anhydrous acid ordinary glacial acetic acid is cooled 
until a considerable mass of crystals is formed. The liquid 
portion, containing more water than the crystals, is poured 
off. By repeating the process a number of times anhydrous 
acetic acid is obtained. 

Glacial acetic acid is very stable. It is not readily decom- 
posed by heat, nor is it oxidized by chromic acid. It is a 
valuable solvent for many organic compounds. 

Sodium Acetate, CH 3 C00Na, is prepared by neutralizing 
dilute acetic acid with sodium carbonate. The solution on 
evaporation yields crystals of the hydrated salt, CH 3 COONa 
+ 3H,0. The crystals lose their water on heating, and the 
anhydrous salt fuses without decomposition at 319°. Sodium 
acetate is used for the preparation of acetyl compounds. 

Potassium Acetate, CH 3 C00K, is a very deliquescent and 
soluble salt. It is used in medicine. 



CH COO 

Calcium Acetate, rjHCOO^ 81 ' * s an easn "y soluble salt, 

crystallizing with two molecules of water. 

CH COO 

Copper Acetate, CH^COO-^ 11 ' * s P re P are ^ ^y dissolving 

copper hydroxide or carbonate in acetic acid. It forms blue 
crystals containing one molecule of water. Verdigris is a basic 
copper acetate formed by the combined action of air and acetic 
acid on copper. It is employed in dyeing. It is very poison- 
ous 



silicon. 363 

Silicon, Si. 

Atomic Weight, 28. Density, 2.49. 

This element occurs in silicon dioxide, Si0 2 , or quartz, and 
in the numerous silicates, which, together with quartz, con- 
stitute most of the rocks, excepting limestones. Silicon is, 
next to oxygen, the most abundant element in the crust of the 
earth. 

Amorphous silicon is obtained by heating a mixture of po- 
tassium silicon fluoride and metallic potassium : 

K 2 SiF 6 + 4K = Si + 6KF. 

The potassium fluoride is dissolved by water, and the silicon 
is left as an amorphous brown powder, which burns readily 
when heated in air to silicon dioxide. 

If the reduction is made with aluminum, and this metal is 
dissolved out of the metallic product by hydrochloric acid, 
silicon will be obtained in dark hexagonal tablets resembling 
graphite. 

Crystalline silicon is also prepared by the following process. 
A mixture of potassium silicon fluoride, zinc, and metallic 
sodium is thrown into a red-hot crucible, and kept hot for 
some time, but not hot enough to volatilize the zinc. After 
cooling, the product is treated successively with hydrochloric, 
hot nitric, and hydrofluoric acids, when dark octahedral 
crystals of silicon will remain. 

Quartz is reduced in the electrical furnace and silicon is 
obtained in a crystalline condition. If the reduction is made 
in the presence of copper this metal dissolves the silicon. 
The addition of a small proportion of silicon to some morals 
and alloys increases their tensile strength. 

Silicon may be fused in a crucible and cast into bars. The 
crystalline form, when heated in oxygen, becomes coated with 
oxide, which protects it from further oxidation. Silicon dis- 



364 THE FOURTH GROUP. 

solves m a boiling solution of potassium hydroxide, with evo- 
lution of hydrogen and formation of potassium'silicate. 

Silicon Tetrahydride, SiH 4 . Gas density, 16. — Pure silicon 
tetralrydride is obtained by treating the compound SiH(OC 2 H 5 ) 3 
with sodium. The colorless gas thus prepared does not ignite 
spontaneously at ordinary temperature, but does so when 
mixed with hydrogen or when slightly warmed. When an 
alloy of silicon and magnesium is treated with hydrochloric 
acid, silicon tetrahydride together with hydrogen is evolved, 
and the gas thus obtained ignites spontaneously on escaping 
into the air, and burns with formation of a cloud of silica. 
Silicon tetrahydride is the analogue of methane, CH 4 . 

Silicon Tetrafluoride, SiF 4 . Gas density, 52. — This com- 
pound is formed when hydrofluoric acid comes into contact 
with silica : 

4HF + Si0 2 = SiF 4 + 2H 2 0. 

The water formed in the reaction decomposes the silicon 
tetrafluoride as stated under Hydrogen Silicon Fluoride. 

Pure silicon tetrafluoride is best prepared by heating a mix- 
ture of fluor spar (calcium fluoride), sand, and sulphuric acid: 

2CaF 2 + 2H 2 S0 4 + Si0 2 = SiF 4 + 2CaS0 4 + 2H 2 0. 

An excess of concentrated sulphuric acid is used to take up 
the water formed. 

Silicon tetrafluoride is a pungent colorless gas, fuming in 
moist air. It combines with dry ammonia to form a white 
compound having the composition SiF..2XH 3 . This sub- 
stance is completely dissociated at high temperatures into sili- 
con tetrafluoride and ammonia. 



HYDROFLUOSILICIC ACID. 365 

Hydrogen Silicon Fluoride or Hydrofluosilicic Acid, H 2 SiF 6 . 

— This acid is formed when silicon fluoride is passed into 
water, silicic acid at the same time separating : 

3SiP 4 + 4H a O = 2H 3 SiF 6 + Si(OH) 4 . 

The saturated solution of hydrofluosilicic acid is a strongly 
acid fuming liquid, which leaves no residue when evaporated 
in a platinum dish, since the acid is decomposed on heating 
into silicon tetrafluoride and hydrofluoric acid. 

A solution of hydrofluosilicic acid is used as a reagent for 
barium, precipitating the metal as barium silicofluoride, 
BaSiF 6 , when added to a solution of a barium salt. 

Exp. 224. — a. Place a mixture of equal parts of fluor spar and quartz 
sand in a flask containing 9 parts of oil of vitriol. Connect the flask 
by means of glass and rubber tubing with a delivery-tube, having a 
diameter of a centimeter or more, and which dips under mercury in a 
cylinder. Heat the flask cautiously with a lamp. Silicon fluoride will 
escape and fume in the air. Pour water into the cylinder. Each 
bubble of gas on entering the water will be surrounded by a film of 
gelatinous silica. After a time the gas will escape through channels 
formed. These should be broken up by stirring. 

Filter the solution of lrydrofluosilicic acid through cloth. 

b. Add some hydrofluosilicic acid to a concentrated solution of bari- 
um chloride. Barium silicon fluoride will be precipitated. 

c. Dry the gelatinous silica, and heat intensely in a crucible. The 
light white powder obtained is pure silica. 

Silicon Tetrachloride, SiCl 4 . Gas density, 84.8.— This sub- 
stance is a colorless fuming liquid, boiling at 59°. C. It is 
formed by the direct union of silicon and chlorine, but is 
best prepared by passing chlorine through a mixture of silica 
and charcoal heated intensely in a porcelain tube: 

SiO a + 20 + 2Cl a = Si01 4 + 2CO. 



366 THE FOUETH GEOUP. 

Silicon tetrachloride is quickly decomposed by water^ "with 
separation of gelatinous silica: 

SiCl 4 + 4H 2 = Si(OH) 4 + 4HC1. 

Silicon Trichloride, Si 2 Cl 6 . Gas density, 134.2.— This com- 
pound is obtained when the tetrachloride is passed over white- 
hot silicon. It is a mobile colorless liquid, solidifying at —1°, 
and boiling at 146°. 

Silicon Chloroform, SiHCl 3 . Gas density, 68.— This body is 
an analogue of chloroform, CHC1 3 . It is formed when silicon 
is heated in dry hydrochloric acid gas. 

Silicon Dioxide or Silica, Si0 2 , occurs in a great variety of 
forms. As quartz it is found in transparent colorless hexagonal 
crystals, terminated with six-sided pyramids. Quartz has a 
hardness of 7, and a density of 2.6. The varieties of native 
silica, on account of their peculiarities, are divided into three 
series, viz., the vitreous variety, having a glassy fracture ; the 
chalcedonic, having a semi-vitreous and waxy lustre, and trans- 
lucent ; the jaspery, having little or no lustre, and opaque. 

Among the vitreous varieties are the pure transparent color- 
less quartz or rock crystal, smoky quartz, common quartz 
sand, the amethyst, rose quartz, and milky quartz, the latter 
occurring in rock masses nearly opaque white. The amethyst 
has a purple or bluish-white color, due to traces of manganese, 
and is valued as a gem. Eose quartz is pink or rose-colored. 
Its color fades on exposure to light. 

The chalcedonic varieties include chalcedony, agate, carne- 
lian, onyx, and flint. 

With the jaspery varieties are classed the various kinds of 
jasper, the bloodstone, touchstone, silicified or petrified wood, 
and some kinds of quartz sandstone. 

Opal is an amorphous form of silica, usually containing a 
little water. It is not quite as hard as quartz. The finer 
varieties are much esteemed as gems. To this variety of 



SILICIC ACIDS. 367 

silica belongs infusorial earth which consists mainly of the 
silicious remains of microscopic plants. It is used for polish- 
ing, and for mixing with nitroglycerine in the manufacture 
of dynamite. 

Silica fuses in the oxyhydrogen flame to a clear glass. It is 
insoluble in water and acids excepting hydrofluoric. Amor- 
phous silica is soluble in solutions of potassium and sodium 
hydroxides and carbonates ; but crystalline silica is almost in- 
soluble in these solutions. When either form is fused with 
an alkali hydroxide or carbonate a soluble alkali silicate is ob- 
tained. 

Silicon Hydroxides or Silicic Acids. — When hydrochloric 
acid is added to a not too dilute solution of an alkali silicate 
a gelatinous mass separates, which is perhaps orthosilicic 
acid, Si(OH) 4 , but which, on drying at ordinary temperatures, 
leaves no well-defined hydroxide or hydrate. If a dilute solu- 
tion of an alkali silicate is poured into hydrochloric acid a 
clear solution results. The alkali salt may be removed by di- 
alysis and a solution obtained containing 5 per cent of silica, 
and which may be concentrated by boiling in a flask until it 
has 14 per cent of silica. The solution on standing changes 
to a transparent jelly. 

Exp. 225. — a. Pour a few cubic centimeters of a solution of sodium 
silicate (commercial water glass will answer) into an excess of hydro- 
chloric acid in a porcelain dish ; mix thoroughly by stirring, and 
evaporate to dryness — best on a water-bath. Moisten the residue with 
hydrochloric acid, add water to dissolve the sodium chloride, and wash 
the silica thoroughly on a filter with hot water. The silica after dry- 
ing contains a little water, which may be driven off by ignition. 

b. Try to dissolve a portion of the silica in hydrochloric acid, and an- 
other portion in a solution of potassium hydroxide. 

c. Dilute sodium silicate with water until the solution does not gelati- 
nize on pouring into hydrochloric acid. Fill a test-tube with the acid 
solution and set aside. If not too dilute, silicic acid will separate as a 
jelly in a few days. Evaporate on a water-bath the remainder of the 
solution of silicic acid until a jelly forms. 



368 



THE FOURTH GROUP. 



Silicates. — The greater number of silicates occur as 
minerals. But few silicates have been prepared in a state 
of purity. The observed gas densities of ethyl orthosilicate, 

^ 0-C.H. 
^^-^"^ r\_ n'xj O— P TT 

Si =zzzZZI O-C FT 5 ' anc *" °^ etn ^ nietasilicate, SiO<Q_p 2 -rr% 
^0-C 2 2 H 5 

correspond to the formulas given. These compounds show 
the tetravalent character of silicon in silicates. Many sili- 
cates may be viewed as derivatives of polysilicic acids formed 
by the dehydration of two or more molecules of orthosilicic 
acid, thus : 



Tetrabasfc 
disilicic acid. 



Dibasic 
disilicic acid. 




Octobasic 
trisilicic acid. 



Hexbasic 
trisilicic acid. 



Tetrabasic 
trisilicic acid. 



Dibasic 
trisilicic acid. 




Si 



Si 



Si 



OH 
:0 


:0 



:o 

-O 

OH 



SILICATES — GLASS. 369 

Sodium Silicates. — "When a mixture of equal molecules of 
silica and sodium carbonate is fused, sodium metasilicate, 
Na 2 Si0 3 , is formed. It is soluble in water, and separates in 
hydrous crystals when the solution is evaporated. If an ex- 
cess of sodium carbonate is used, sodium trisilicate, Na 8 Si.O 10 , 
is obtained. Commercial silicate of soda, known also as water 
glass and soluble glass, is a mixture of sodium silicates, con- 
sisting chiefly of sodium tetrasilicate, Na 2 Si 4 9 . It is manu- 
factured by heating together 180 parts of sand, 100 of soda- 
ash, and 3 of charcoal. The glassy mass which results dis- 
solves slowly in boiling water, forming a viscid alkaline 
liquid. Soluble glass is used with fresco colors, as a cement 
for mending broken porcelain and stone-ware, and in large 
quantities in the manufacture of silicated soaps. 

Potassium Silicates resemble the sodium silicates, but are 
less used on account of the high cost of the potash required 
for their manufacture. 

Glass is valued chiefly for its transparency and durability. 
There is no substitute for window glass. It protects from the 
weather and admits sunlight,, which not only renders dwell- 
ings cheerful, but acts also as a purifying agent. Light, too, 
makes filth conspicuous. Glass was used by the ancients, but 
cheap glass for the multitude came with the great development 
of the chemical industries in the present century. 

Glass is a mixture of silicates of metals belonging to two 
different groups. The ordinary kinds of glass are composed 
of alkali silicates and calcium or lead silicate. The alkali 
silicates are glassy in appearance after fusion, but are soluble 
in water. The calcium and lead silicates are practically 
opaque, and are decomposed by acids. But when sodium 
silicate and calcium silicate are melted together, a glass is 
formed which is not soluble in water, and is not corroded by 
acids, excepting hydrofluoric acid. 



370 THE FOURTH GEOUP. 

Window or crown glass is made by melting in a fire-clay pot 
a mixture of sand, soda-ash, and lime or limestone. Broken 
glass (cullet) is also added to facilitate the melting. Arsenic 
trioxide and manganese dioxide are used in small quantities 
to counteract color imparted by iron contained in the ingre- 
dieuts. Window glass consists of 70 to 75 per cent of Si0 2 , 
and from 12^ to 15 per cent each of Xa 2 and CaO. 

Flint or crystal glass is composed of potassium and lead sili- 
cates. It is characterized by its high density, brilliant lustre, 
and great refracting power. It is more fusible than other 
kinds of glass, and does not resist the action of chemicals as 
well. It is used for the finer kinds of glass-ware, and for op- 
tical purposes. 

Bohemian or hard glass is a potassium calcium silicate. It 
is valuable for its infusibility and resistance to chemicals. 

Common green bottle glass is similar in composition to 
window glass, but contains aluminum and iron derived from 
the impure and cheap materials used in its manufacture. The 
green color of bottle glass is due to iron. 

The brittleness of glass, and its property of gradually 
softening and becoming pasty when heated, are familiar to 
chemical students, and need not be described. Glass is slowly 
acted upon by water ; some saline solutions dissolve it to a 
greater extent, and alkaline solutions corrode it still more. 
Hence many exact chemical experiments must be made in 
platinum instead of glass dishes. 

Water from steam condensed in glass will leave more residue 
on evaporation than if condensed in a block-tin or platinum 
tube. 

Clay. — Kaolin, the purest kind of clay, is a hydrous alumi- 
num silicate. It has been formed by the decomposition of 
aluminous minerals, especially the feldspars, which are sili- 
cates of aluminum and sodium or potassium, often containiug 
small amounts of other metals. Common clays are composed 






PORCELAIN — STON-E WARE — BRICK — CEMENT, 371 

of kaolin mixed with powdered feldspar, quartz, or other 
minerals. Clays are distinguished by forming a plastic mass 
with water, which becomes hard and compact when intensely 
heated. The properties of different clays will be noticed in 
connection with articles manufactured from them. 

Porcelain is made by mixing fine white kaolin with sufficient 
pulverized quartz and feldspar so that the mass will partly fuse 
when heated and become translucent. The ware is next 
coated with a more fusible mixture, which on heating forms 
the smooth, glass-like surface of porcelain. 

Common Earthen and Stone Ware are made from the com- 
mon varieties of clay. Both kinds of ware before glazing are 
porous and opaque. 

Bricks are of two kinds, common building brick and fire- 
brick. The former are often red from oxide of iron con- 
tained in the clay. Fire-brick are made by mixing fire-clay 
with pulverized and previously burned fire-clay to prevent 
cracking on drying. Fire-brick are used for stove and fur- 
nace linings. The crucibles used in metallurgical processes, 
and the pots in which glass is melted, are made of the most 
refractory fire-clay. 

Hydraulic Mortar or Cement. — Limestone containing 10 per 
cent or more of clay is converted by burning into a mass, 
which in state of fine powder possesses the property of harden- 
ing under water. The chemical change which occurs in the 
hardening is not well understood. It has been supposed that 
an insoluble hydrous calcium aluminum silicate is formed, 
and also that the calcium aluminate contained in the cement 
takes up water. Some cements, containing little silica and 
alumina, require weeks to harden ; others made from hydraulic 
limestone containing 25 to 35 per cent of clay begin to harden 



372 THE FOURTH GROUP. 

soon after wetting, and become firm in a few hours. Cement 
is obtained not only from hydraulic limestones, but also by 
calcining a mixture of chalk or limestone and clay. 



Titanium, Ti. 

Atomic Weight, 48. 

Titanium occurs in nature as dioxide, and is also found in some iron 
ores and silicates. It is a comparatively rare element. Metallic tita- 
nium is obtained by reducing potassium titanium fluoride with potas- 
sium. On dissolving the fused mass in water the titanium remains as 
an amorphous powder, which burns brilliantly in air, is soluble in acids, 
and decomposes boiling water. According to Kern, the latter reaction 
is due to potassium contained in the product. Crystals of titanium are 
formed when vapor of titanium chloride is passed over molten sodium. 

Titanium Dichloride, TiCl 2 , is formed by heating the higher chlorides 
in a current of hydrogen. It is a hygroscopic brown powder, which 
decomposes water. 

Titanium Trichloride, TiCl 3 , is obtained in dark-violet scales when the 
vapor of titanium tetrachloride and hydrogen are passed together 
through a red-hot tube. It is non-volatile, and decomposes at high tem- 
peratures into TiCl 2 and TiCl 4 . It is deliquescent, and yields a violet 
solution with water. 

A solution prepared by dissolving metallic titanium in hydrochloric 
acid deposits on evaporation crystals of the hydrated trichloride, TiCl 3 
-\- 4H 2 0. During the evaporation titanic acid separates, and the solution 
must be frequently filtered. 

Titanium trichloride is a strong reducing agent, separating sulphur 
from sulphur dioxide. 

Titanium Tetrachloride, TiCl 4 , is formed by heating titanium in chlo- 
rine, and by passing chlorine over a mixture of titanic oxide and char- 
coal. It is a colorless liquid, boiliDg at 136°, having a gas density cor- 
responding to the formula TiCl 4 . 

Titanium tetrachloride is soluble in cold water, and forms in moist 
air a soluble hydrate. The aqueous solution, on heating and evaporat- 
ing, decomposes into titanic and hydrochloric acids. 



TITANIUM. 373 

Titanium Tetrafluoride, TiF 4 . — When a mixture of fluor spar, titanic 
oxide, and fuming sulphuric acid is distilled, titanium tetrafluoride is 
obtained as a colorless fuming liquid. 

Hydrogen Titanium Fluoride, H 2 TiF 6 , is formed when titanic oxide is 
dissolved in hydrofluoric acid. 

Potassium Titanium Fluoride, K 2 TiF 6 , is prepared by neutralizing hy- 
drogen titanium fluoride with potassium hydroxide, and also by fusing 
titanic oxide with potassium carbonate and dissolving the fused mass 
in boiling dilute hydrofluoric acid. On cooling, the potassium salt 
separates in scales which melt without decomposition. 

Titanium Sesquioxide, Ti 2 3 , is prepared by igniting the dioxide in a 
current of hydrogen. It is soluble in sulphuric acid, forming a violet 
solution. 

Titanium Dioxide or Titanic Oxide, Ti0 2 , occurs as rutile in tetragonal 
crystals, density 4.18-4.25; as octahedrite (anatase) ; also in tetragonal 
crystals of totally different form, having a density of 3.82-3.95 ; and as 
brookite in orthorhombic crystals, density 4.12-4.23. Amorphous tita- 
nium dioxide is prepared by igniting titanium hydroxide. It is a white 
powder, insoluble in water, hydrochloric and dilute sulphuric acids, 
but dissolves in hydrofluoric acid and concentrated sulphuric acid. 

The amorphous dioxide is converted into crystals of octahedrite at 
about 860° ; into crystals of brookite at about 1000°; and at still higher 
temperatures into crystals of rutile. 

Titanic Acids. — Titanium forms two hydroxides, which are called 
titanic and metatitanic acids. According to some authorities, both acids 
have the composition Ti(OH) 4 , while others assign the formula 
TiO(OH) 2 . Both acids lose water with increasing temperature, and 
finally leave Ti0 2 . 

Titanic acid is obtained by adding an alkali hydroxide to a cold solu- 
tion of titanic chloride. It is soluble in hydrochloric, dilute sulphuric, 
and nitric acids, and the solution on boiling deposits metatitanic acid, 
which is insoluble in acids except hydrofluoric and concentrated sul- 
phuric. Metatitanic acid is also obtained by dissolving titanic oxide 
in sulphuric acid, diluting the solution Avith a large quantity of water, 
and then boiling. 

The titanium hydroxides act as weak bases, and also as acids forming 
salts both with acids and bases, but the compounds are not well under- 
stood. The following formulas illustrate the complex composition of 
titanatcs: K 2 Ti() 3 + 4II,(), K 2 Ti 3 7 + 2H a O, K 9 Ti 6 ls + 2H a O, and 
K 2 Ti 3 7 -}-3H 2 0. The iron titanatcs are even more complex. 

Titanium Sesquisulphate or Titanous Sulphate, Tij(S0 4 ) ;1 -f- 8H 2 0, is 



374 THE FOURTH GROUP. 

formed by dissolving titanium in sulphuric acid. The violet solution 
on concentration becomes blue, aud deposits crystals of the salt. 

Titanium Eisulphate or Titanic Sulphate, Ti(S0 4 ) 2 + 3H 2 0, is obtained 
as a yellow amorphous mass when titanous sulphate is treated with 
nitric acid and a few drops of sulphuric acid, and the solution is evapo- 
rated. 

Titanyl Sulphate, TiOS0 4 , is obtained by dissolving titanic oxide in 
boiling sulphuric acid and evaporating the solution. 

Titanium Nitrides.— The compound Ti 3 N 4 results when titanium tetra- 
chloride and ammonium chloride are heated together. By passing 
dry ammonia over red-hot titanic oxide the mono nitride TIST is formed. 

Titanium Cyano-nitride, Ti(CN) 2 .3Ti 3 N 2 , is formed in blast furnaces 
in the smelting of titaniferous iron ores, and is found in the slag and 
pig iron as copper-colored scales. It is also formed when titanic oxide 
and charcoal are intensely heated in a crucible, the nitrogen being sup- 
plied by the air 



Zirconium^ Zr. 

Atomic Weight, 90.7. Density, 4.15. 

The oxide of zirconium was discovered by Klaproth in 1789 in the 
mineral zircon, ZrSi0 4 . The metal was first isolated by Berzelius, 
who obtained it in the form of an amorphous dark powder by reducing 
potassium zirconium fluoride with potassium. Crystalline gray scales 
of the metal are obtained when the reduction is made with aluminum. 
Crystalline zirconium is brittle, very hard, and infusible. It does not 
oxidize at a red heat, but at a white heat becomes coated with oxide, 
which protects it from further oxidation. It is soluble in aqua regia 
and in hydrofluoric acid. 

Zirconium Chloride, ZrCl 4 , is prepared by heating an intimate mixture 
of zirconium oxide and charcoal in a current of chlorine. It forms a 
white sublimate. Its observed gas density is 117; theory for ZrCl 4 is 
115.5. Its aqueous solution loses hydrochloric acid when evaporated, 
and deposits hydrous crystals of zirconyl chloride, ZrOCl 2 . The same 
compound is also formed when zirconium hydroxide is dissolved in 
hydrochloric acid. 

Zirconium Fluoride, ZrF 4 , is obtained \>y heating zirconium oxide with 
acid ammonium fluoride. A solution of the fluoride in hydrofluoric 
acid yields crystals of ZrF 4 + 3H 2 0, which decompose on drying, and 
at a red heat leave zirconium oxide. 



ZIRCONIUM — CERIUM. 375 

Potassium Zirconium Fluoride, K 2 ZrF B , is prepared by fusing zircon 
with acid potassium fluoride, or by adding potassium fluoride to an 
excess of a solution of zirconium fluoride. It is very soluble in hot and 
sparingly soluble in cold water. 

Zirconium Oxide or Zirconia, Zr0 2 , is obtained as a white powder by 
igniting the hydroxide. When zirconia is intensely heated with borax, 
crystals are formed which are isomorphous with cassiterite and rutile. 

Zirconium Hydroxide, Zr(0H) 4 , separates as a gelatinous white precipi- 
tate on addition of ammonia to a cold solution of the sulphate or chlo- 
ride. The washed precipitate after drying at 17° has the composition 
Zr(OH) 4 (Hermann). Berzelius found that it corresponded toZrO(OH) 2 , 
probably after drying at a higher temperature. Zirconium hydroxide 
when precipitated cold is easily soluble in hydrochloric acid, but when 
precipitated hot is difficultly soluble. It is insoluble in alkali hydrox- 
ides. Zirconium hydroxide is both acidic and basic, as the zirconates 
and zirconium sulphate show. 

Zirconates. — Sodium metazirconate, ^rO<Q~^^, is formed when a 

mixture of zirconia and sodium carbonate is fused. If the fusion 
is kept at a white heat for some time sodium orthozirconate, 
,0-Na 

Zr< 0-Na' isfarmed ' 

^O-Na 

Zirconium Sulphate, Zr(S0 4 ) 2 , is obtained by evaporating a solution of 
the oxide or hydroxide in sulphuric acid, and heating the residue nearly 
to redness. A solution of the sulphate saturated with zirconium hydrox- 
ide, yields on evaporation zirconyl sulphate, ZrOS0 4 . 



Cerium, Ce. 

Atomic Weight, 141. Density, Q.6. 

This element was discovered in 1803 by Klaproth, and independently 
by Berzelius and Hisinger. It occurs associated with lanthanum and 
didymium m a few rare minerals. The metal is obtained by elect ro- 
lyzing the chloride. It has the color and lustre of iron, is tolerably 
permanent in dry air, and in moist air becomes first yellow, then blue. 
and finally green. Its melting point is below that of silver. Cerium 
burns in air with greater brilliancy than magnesium. It is soluble in 
acids. 



376 THE FOURTH GROUP. 

Cerium forms two classes of compounds, viz., the cerous, in which 
cerium is trivalent; and the eerie, in which it is tetravalent. 

Cerous Chloride, CeCl 3 , is formed when the metal is heated in chlorine, 
A solution of cerous oxide in hydrochloric acid on evaporation deposits 
crystals of CeCl 3 + 5H 2 0, which on heating give off hydrochloric acid 
and change to a basic salt. 

Cerous Oxide, Ce 2 3 , is obtained by igniting the carbonate or oxalate 
in pure hydrogen. It is a bluish powder, which quickly oxidizes in air 
to dioxide. Sodium hydroxide precipitates from solutions of cerous 
salts a hydroxide which absorbs oxygen and carbon dioxide from the air. 

Cerous Sulphate, Ce 2 (S0 4 ) 3 , is prepared by dissolving cerous oxide or 
carbonate in sulphuric acid. It is very soluble in cold and only slightly 
soluble in hot water. This phenomenon is explained by the existence 
of different hydrates of the salt which differ in solubility. 

Cerous Carbonate, Ce 2 (C0 3 ) 3 + 9H 2 0, forms a bulky precipitate when 
ammonium carbonate is added to a solution of cerous sulphate. 

Ceric Oxide or Cerium Dioxide, Ce0 2 , is formed when a cerium salt of a 
volatile acid is ignited in air. It dissolves in hydrochloric acid, and 
the solution on heating evolves chlorine, and cerous chloride is formed. 

Ceric Sulphate, Ce(S0 4 ) 2 , has been obtained in hydrous crystals. A 
solution of ceric oxide in sulphuric acid evolves ozone, and contains 
cerous sulphate and ceric sulphate. 



Thorium, Tli. 

Atomic Weight, 232. Density, 11. 

The oxide of this metal, thoria, was discovered in 1829 by Berzelius 
in a mineral now known as thorite, a thorium silicate. Thorium occurs 
in a few other rare minerals. Nilson isolated the metal by the follow- 
ing process. A mixture of dry potassium thorium chloride, sodium 
chloride, and metallic sodium was heated to redness in an iron crucible 
closed by a cover held \>y a clamp. The contents of the crucible after 
cooling were treated with water to dissolve the salts, and also the metal- 
lic sodium remaining. The thorium which remained was washed with 
alcohol and ether, and dried at 100°. The metal thus obtained was in 
the form of a gray crystalline, brittle powder, which ignited in air 
below redness and burned brilliantly to a snow-white oxide. It has 
not been fused. It is soluble in concentrated hydrochloric acid, but is 
not attacked by solutions of alkali hydroxides. 



THORIUM — GERMANIUM. 377 

Thorium Chloride, ThCl 4 , is prepared by heating the metal in chlorine, 
and also in hydrogen chloride. Vapor density found, 178; theory for 
ThCU, 187. A solution of the chloride is obtained by dissolving the 
hydroxide in hydrochloric acid. 

Thorium Dioxide, Th0 2 , remains as an amorphous powder when the 
hydroxide is ignited. On heating the powder intensely with borax, it 
changes to crystals which are isomorphous with crystals of the com- 
pounds Sn0 2 , Zr0 2 . and Ti0 2 . 

Thorium Hydroxide, Th(OH) 4 , separates on addition of ammonia to a 
solution of thorium chloride. It absorbs, when moist, carbon dioxide. 
It does not appear to possess acidic properties. 

Thorium Sulphate, Th(S0 4 ) 2 , is prepared by dissolving the oxide in hot 
sulphuric acid, and then expelling the excess of acid by heat. It is 
soluble in water, and forms several hydrates. 



Germanium, (*e. 

Atomic Weight, 72. Density, 5.5. 

This element was discovered in 1886 by Winkler in a mineral found 
in the mines at. Freiberg, in Saxony. Metallic germanium is obtained 
by heating the oxide in hydrogen. It melts at about 900°, has a gray- 
ish-white color, exhibits a beautiful metallic lustre, and crystallizes in 
regular octahedrons. It is insoluble in hydrochloric acid, soluble in 
aqua regia, and is converted into oxide by nitric acid. It is a tetrad, 
forming, however, compounds in which it is a dyad. 

Germanium Dichloride, GeCl 2 , is probably formed when the monosul- 
phide is heated in hydrogen chloride. It has not been obtained in the 
pure state. 

Germanium Tetrachloride, GeCl 4 , results from the direct combination of 
its elements. It is a colorless liquid, boiling at 86°. At —100° it does 
not solidify. Gas density found, 107.4; calculated for GeCl 4 . 106.8. 

Germanium Chloroform, GeHCl 3 , is prepared by healing the metal in 
hydrogen chloride. It is a colorless mobile liquid, which absorbs 
oxygen from the air with formation of oxychloride, GeOCU(?), and 
evolution of hydrogen chloride. 

Germanium Tetrafluoride, GeF 4 . — Germanium oxide dissolves in hydro- 
fluoric acid, and the solution yields on evaporation over sulphuric acid 
crystals of GcF 4 + 31I 2 0. When the vapors from a heated mixture of 
calcium fluoride and an excess of germanium oxide and sulphuric acid 



378 THE FOUKTH GEOUP. 

are conducted into water, a strongly acid solution results, which is sup- 
posed to contain hydrogen germanium fluoride, H 2 GeF 6 . 

Potassium Germanium Fluoride, K 2 GeF 6 , separates on mixing solutions 
of germanium fluoride and potassium fluoride. It is sparingly soluble 
in cold and more soluble in hot water. 

Germanium Monoxide, GeO, results from the decomposition of the di- 
chloride with potassium hydroxide. 

Germanium Dioxide, Ge0 2 , is formed when the metal is burned in air, 
or is oxidized by nitric acid. It is best prepared by decomposing the 
tetrachloride with water and igniting the precipitate. It is somewhat 
soluble in water. Salts of the dioxide appear to exist, but have not 
been obtained of definite composition. 

Germanium Monosulphide, GeS, is obtained by heating the disulphide 
for some time in hydrogen. Its observed gas density at 1100° is 51, 
theory requiring 52. 

Germanium Disulphide, GS 2 , separates as a white precipitate when a 
solution of the dioxide in sulphuric acid is treated with hydrogen sul- 
phide. It is soluble in 222 parts of water. 



Tin (Staimum)j Sn. 

Atomic Weight, 118. Density, 7.3. 

Tin was an article of commerce before the Christian era. 
It is said to have been found native in small quantities. It 
occurs chiefly as dioxide, known as tin stone. There are but 
few tin mines, and the metal is not abuudant. The tin stone 
is separated from the gangue by mechanical and. chemical 
processes, and is reduced by heating with coal. The metal is 
purified, if necessary, by liquation. In this process the crude 
tin is gradually heated in a furuace, the pure tin flows off 
first, leaving less fusible alloys of tin. Commercial tin is 
liable to coutain lead, copper, iron, zinc, antimony, and 
arsenic. Tin from Banca is fairly pure, containing only traces 
of impurities. 

Tin is almost as white as silver, and is harder than lead, 
and somewhat softer than gold. It has a crystalline structure, 



tin. 379 

which causes the crackling or " tin cry" when a bar of the 
metal is bent. It is very malleable. At 100° it may be drawn 
into wire, but at 200° it is so brittle that it may be pulver- 
ized. At low temperatures ( — 40°) tin slowly becomes granu- 
lar, and falls to a powder. It melts at 230°, undergoing 
slight oxidation in air, and at higher temperatures it burns to 
dioxide. At ordinary temperatures the lustre of tin is but 
little impaired by exposure to air and water. Tin dissolves 
in hydrochloric and dilute nitric acid, and is attacked by hot 
concentrated solutions of potassium and sodium hydroxides. 

Since tin does not tarnish in air, and is not corroded by 
vinegar or other liquids used in cookery, the metal is especially 
valuable for culinary vessels. It is not sufficiently rigid for 
ordinary articles unless these are made quite heavy, and is, 
moreover, rather costly. Hence the metal is largely used as 
a protecting coating for iron and copper. The process of 
tinning sheet copper is as follows. The surface of the copper 
is carefully cleaned, wet with hydrochloric acid, to prevent 
formation of any oxide, and then molten tin is poured over 
the sheet, the excess of the tin being wiped off with cotton 
waste. Sheet iron is tinned by immersion in a bath of molten 
tin. After removal from the bath it is placed in hot tallow, 
so that the excess of tin may flow off. In order that the tin 
may adhere to the iron the latter must be carefully cleaned 
and protected from the air by a coating of tallow. 

In place of pure tin an alloy of tin and lead, which is less 
costly, is largely used for tinning iron and copper. Such a 
coating has a bluish-leaden lustre. It answers for some pur- 
poses, but is quickly corroded by ordinary water, and is totally 
unfit for culinary vessels on account of the lead. 

Tin may be deposited by electrolysis. When only a thin 
coating of tin is desired, brass and copper articles are tinned 
by immersion in a solution of the metal in which pieces of 
metallic tin are placed. Mirrors are coated with an amalgam 
of tin. 



380 THE FOURTH GROUP. 






Tin is used in the form of foil for various purposes, but 
most of the so-called tin foil is chiefly lead between two sur- 
faces of tin, and is made by rolling out plates of lead coated 
with tin. 

Tin forms two classes of compounds, viz., the stannous, con- 
tainiug bivalent tin ; and the stannic, in which tin is tetrava- 
lent. Stannous tin is always basic, but stannic tin is acidic or 
basic according as it is combined with basic or acidic radi- 
cals. 



Stannous Compounds. 

CI 
Stannous Chloride, Tin Dichloride, SnCl 2 or Sn< cl , is ob- 
tained by heating tin in dry hydrogen chloride or with mer- 
curic chloride. Its gas density at 900° corresponds to that of 
molecules of SnCl 2 . Tin dissolves in hydrochloric acid, and 
the solution on evaporating and cooling deposits crystals of 
the hydrate SnCl 2 -f 2H 2 0, known as tin salt. It dissolves 
in 0.37 part of water, but is decomposed by a large quantity 
of water with formation of a basic salt having the composition 
2SnOHCl -p H 2 0. This basic salt is insoluble in water, but 
soluble in hydrochloric acid. Tin salt, both in the dry state 
and in solution, absorbs oxygen from the air with formation 
of a white insoluble oxychloride. Stannous chloride is a 
strong reducing agent, separating metallic gold from a solu- 
tion of gold chloride, and reducing cupric chloride to cuprous 
chloride. In these reactions stannous chloride is changed to 
stannic chloride. Tin salt is extensively used as a mordant in 
dyeing. 

Stannous Oxide, SnO or Sn=0. — Potassium carbonate pro- 
duces in a solution of stannous chloride a white precipitate of 
Sn< OH 

stannous oxy hydroxide, Q . This when heated in carbon 
kn< 0H 



STANNIC COMPOUNDS. 381 

dioxide is converted into stannous oxide. The oxide takes 
fire when touched with a glowing splinter, and burns 
brilliantly, stannic oxide being formed. 

Stannous Sulphide, SnS, results when tin foil burns in sul- 
phur vapor. It forms a green vapor at high temperatures. 
Hydrogen sulphide throws down from stannous solutions a 
dark-brown precipitate of hydrous stannous sulphide. 

Stannous Nitrate is obtained by dissolving tin in cold dilute 
nitric acid, part of the acid at the same time being reduced 
to ammonia : 

4Sn+10NO,-OH = 4™^>Sn + N0 2 -0-NH 4 + 3H 2 0. 

A solution of stannous nitrate decomposes on heating with 
separation of stannic hydroxide. 

Stannous Sulphate, SnS0 4 or S0 2 < Q > Sn. — This salt is formed 

when concentrated sulphuric acid is heated with an excess of 
tin. Stannous sulphate is readily soluble in water, but the 
solution on standing deposits a basic salt. 



Stannic Compounds. 

/CI 
Stannic Chloride, ) ~ m ^ C „^1-C1 m,. 
Tin Tetrachloride, [ SnC1 * or Sn cr - This com- 

pel 
pound is obtained by passing dry chlorine into molten tin in 
a retort ; the vapor of the chloride is condensed in a cooled 
receiver. Stannic chloride is a colorless mobile liquid, boiling 
at 120°, and having a gas density corresponding to the for- 
mula SnOl,. Exposed to moist air it takes up three molecules 
of water, and when its aqueous solution is evaporated at a 
moderate temperature crystals of the compound Sn01 4 +5H 9 
are deposited. A solution of stannic chloride, much used in 



382 THE FOURTH GROUP. 

dyeing, is prepared by dissolving tin in aqna regia, kept cool 
to avoid the formation of metastannic acid. Dilute aqueous 
solutions of stannic chloride decompose on standing or when 
warmed, with separation of stannic hydroxide. Stannic 
chloride absorbs ammonia with formation of the compound 
SnCl 4 .2NH 3 . This is dissociated by heat into stannic chloride 
and ammonia gases, which recombine on cooling. 

Hydrogen Stannic Chloride or Chlorostannic Acid, H 2 SnCl 6 . 

— This compound is prepared by dissolving tin tetrachloride 
in the required amount of concentrated hydrochloric acid, and 
then passing in hydrochloric acid gas as long as it is absorbed. 
On cooling, crystals separate having the composition H 2 SnCl 6 
+ 6H 2 0. The ammonium and other salts of this acid have 
long been known, but the acid has only recently been iso- 
lated. 

Ammonium Stannic Chloride, (NH 4 ) 2 SnCl 6 , is prepared by 
mixing concentrated solutions of ammonium chloride and 
stannic chloride, when a white crystalline meal separates. 
The salt crystallizes in octahedrons, is permanent in air, and 
is not decomposed in concentrated solution by boiling. A 
dilute solution deposits stannic hydroxide on heating. This 
change renders the salt valuable in dyeing, since the separated 
hydroxide acts as a mordant. 

Stannic Oxide, Tin Dioxide, Sn0 2 or 0=Sn=0, occurs 
native as cassiterite or tin stone. It is formed artificially by 
heating tin in air or by igniting stannic hydroxide. The 
amorphous stannic oxide thus obtained may be changed into 
microscopic crystals of the same form as cassiterite by heating 
in hydrochloric acid gas. Stannic oxide is fusible only at 
a very high temperature, and is insoluble in acids, excepting 
concentrated sulphuric acid, with which it forms a syrupy 
liquid, the dioxide separating again on dilution with water. 
Stannic oxide is rendered soluble by fusion w T ith potassium or 



STANNIC COMPOUNDS. 383 

sodium hydroxide, and with a mixture of sodium carbonate 
and sulphur. It is easily reduced to the metal when heated 
in contact with charcoal or reducing gases. It is used in the 
manufacture of opaque white glass. 

Stannic Hydroxides exist in two modifications, known as 
stannic and metastannic acids, differing in their behavior 
towards solvents, and forming two classes of salts. 

Stannic Acid is formed as a white precipitate when a solu- 
tion of stannic chloride is treated with calcium carbonate not 
in excess, and is also obtained when an acid is carefully added 
to a solution of a stannate. It is slightly soluble in water, and 
reacts acid. Dried over sulphuric acid it has the composition 
H 2 Sn0 3 . It dissolves easily in hydrochloric, nitric, and sul- 
phuric acids, and in solutions of the caustic alkalies. 

Potassium Stannate, K 2 Sn0 3 , is formed when stannic oxide 
is fused with potassium hydroxide. It is readily soluble in 
water, and crystallizes as K 2 Sn0 3 -[- 3H 2 0. 

Sodium Stannate, Na 2 Sn0 3 , is prepared in the same way as 
the corresponding potassium salt, and also by heating tin with 
sodium hydroxide and sodium nitrate. The solution yields 
crystals of Na 2 Sn0 3 +3H 2 0. The salt is used in calico printing. 

The alkali stannates are the only soluble stannates, aud 
from their solutions metallic stannates are precipitated on the 
addition of solutions of metallic salts. 

Metastannic Acid. — Tin and concentrated nitric acid react 
violently with formation of a white insoluble powder, which 
after drying at 100° has the composition H.,Sn0 3 . The salts 
of metastannic acid indicate that it has the formula H 10 Sn 6 O lt . 
Metastannic acid, dissolves, slowly in solutions of potassium and 
sodium hydroxide with formation of potassium and sodium 
metastannates. 



384 THE FOURTH GEOUP. 

Sodium Metastannate, Na 2 H 8 Sn 5 15 , is a slightly soluble crys- 
talline powder obtained by treating metastannic acid with a 
cold solution of sodium hydroxide. 

Stannic Sulphide, SnS 2 , is not formed by the direct union of 
its elements. It may be obtained by heating a mixture of tin 
amalgam, sulphur, and ammonium chloride in a retort. Part of 
the stannic sulphide will remain in the bottom of the retort, 
and part will sublime in golden-yellow crystals. The crystal- 
line stannic sulphide is used as a pigment under the name of 
mosaic gold. Hydrogen sulphide in excess produces in solu- 
tions of stannic chloride a yellow precipitate of stannic sulphide 
which is soluble in alkali sulphides, with formation of sulpho- 
stannates. The compounds Na 2 SnS 3 + H 2 and Na 4 SnS 4 -f- 
12H 2 have been prepared. The former yields, on treat- 
ment with hydrochloric acid, sulphostannic acid, H 2 SnS 3 . 

Exp. 226.— Granulate tin by pouring the molten metal in a thin stream 
into water. 

Exp. 227. — Moisten a clean sheet of copper with hydrochloric acid, 
then melt a bit of tin on the copper, and rub the tin over the copper with 
a rag moistened with acid. 

Exp 228. — a. Digest pure tin foil or granulated tin with concentrated 
hydrochloric acid for some time. To half of the solution add potassium 
chlorate until chlorine is evolved. Expel the free chlorine by heating 
the solution. The potassium chlorate will convert the stannous into 
stannic chloride, the chlorate being reduced to chloride. 

b. Try the reaction of hydrogen sulphide with dilute solutions of 
stannous chloride, and of stannic chloride. Dissolve the moderately 
washed precipitate of stannous sulphide in the least possible quantity 
of yellow ammonium sulphide. The solution when aeidiiied with hy- 
drochloric acid wull give a yellow precipitate, showing that the tin had 
changed from the stannous to the stannic state. 

Exp. 229. — Treat tin with concentrated nitric acid. What compound 
of tin is formed ? 



LEAD. 385 

Lead (Plumbum), Pb. 

Atomic Weight, 207. Density, 11.35. 

Metallic lead occurs very sparingly in nature, but its com- 
pounds are abundant and widely distributed. The most im- 
portant ore of lead is galena, PbS. The smelting of galena 
when free from sulphides of other metals is very simple. The 
ore is roasted in a suitable furnace until the sulphur is burned 
away and lead remains. At first a portion of the lead sul- 
phide is converted by the oxygen of the air into lead oxide and 
sulphate. The heat is then raided and the oxide and sulphate 
react with undecomposed sulphide to form sulphur dioxide, 
which passes off, leaving metallic lead: 

2PbO +PbS = 3Pb + S0 2 ; 
PbS0 4 -f PbS = 2Pb + 2S0 2 . 

Lead has a light bluish-gray color, and the fresh-cut surface 
shows high lustre, which is soon dimmed by superficial oxida- 
tion. It is soft, easily scratched and cut, and leaves a bluish 
mark on paper. It ranks sixth among metals in malleability, 
and eighth in ductility. It has little tensile strength, a wire 
T * ¥ of an inch in diameter breaking with a load of 20 pounds. 
It melts at 334°, and volatilizes between 1400° and 1600°. 
On solidifying it contracts to such a degree that it is poorly 
adapted for making castings. Molten lead oxidizes in the air, 
and when the temperature is sufficient to melt the oxide, the 
latter is absorbed to a slight extent by the lead, rendering the 
metal harder. The lead of commerce is sometimes very pure, 
containing only slight traces of copper, iron, and silver. It is 
also liable to contain antimony, zinc, nickel, and bismuth. 

Lead is only slightly attacked by hydrochloric acid and 
dilute sulphuric acid, but is dissolved by warm concentrated 
sulphuric acid. The best solvent for lead is moderately strong 



386 THE FOURTH GROUP. 

hot nitric acid. Soluble lead salts are poisonous, and when 
taken continuously into the system, even in small quantities 
produce nervous prostration, paralysis, and other disorders. 
Large doses of the soluble salts, as lead acetate, produce 
acute poisoning, and may cause death in a few days. Metallic 
lead is not poisonous to handle, and hence plumbers are not 
subject to lead poisoning. 

The action of drinking waters upon lead is of interest, since 
water which passes through lead pipes may be contaminated 
with lead. 

Lead is not corroded by dry air nor by water free from air, 
while in water containing air it tarnishes quickly. If the 
water is soft, such as rain water or ordinary distilled water, a 
white coating of lead hydroxide is rapidly formed, which is 
converted into a basic carbonate by the carbon dioxide ab- 
sorbed from the air. The white compounds formed by the 
combined action of air and water on lead do not adhere to the 
metal, but scale off, thus exposing fresh surfaces and allowing 
the corrosion to continue. When the surface of lead is large 
in proportion to the bulk of water the latter soon becomes 
milky white from the suspended particles of lead carbonate 
and hydroxide, only a small quantity of the latter being dis- 
solved by the water. 

Drinking waters, rain water excepted, often contain salts 
which have a marked influence on the corrosion of lead. 
Ammonium nitrate, which is one of the products of the 
decomposition of organic matter in water, increases the action 
on the lead, while other nitrates appear to have little effect. 
Hard water is less liable to take up lead than soft water, owing 
to formation of an insoluble and coherent scale by the 
sulphates and carbonates of hard water, the scale preventing 
further corrosion. As a rule, water which has stood in lead 
pipes or in lead or pewter vessels is unfit for drinking, but 
water is not contaminated by passing rapidly through a lead 
pipe. 



LEAD. 387 

Exp. 230. — a. Place a bright strip of lead in a bottle of distilled water. 
b. Place another piece of lead in the drinking water at hand. Allow 
the bottles to stand for some time, best loosely stoppered. The corro- 
sion of the lead in the distilled water will continue for years if prepara- 
tion a is kept, while the action of the drinking water will depend upon 
the impurities it contains. 

Alloys of Lead and Tin. — These metals mix on fusion in 
all proportions. The following table gives the composition 
and melting points of some of the alloys : 

Melting Point. 

Tin, ..... 230° 

Lead, . 334° 

No. 1, 1 2 226° 

" 2, 1 1 188° 

" 3, 2 1 170° 

"4, 3 1 178° 

The solder used for soldering tinware and joining lead pipes 
in plumbing varies much in composition. No. 1 is coarse, 2 is 
common, and 3 is fine solder. Solder is not only more fusible 
than either of the metals, but is much harder and stronger. 

Lead in most of its combinations is bivalent, its tetravalent 
character only appearing in lead methyl, Pb(CH 3 ) 4 , lead ethyl, 
Pb(C 2 H 5 ) 4 , and a few other compounds. 

CI 
Lead Chloride, PbCl 2 , or Pb< cl , is obtained by treating lead 

oxide, carbonate, or sulphide with hydrochloric acid, and also 
as a white crystalline precipitate on addition of a soluble 
chloride to a solution of a lead salt. It is soluble in 135 parts 
of water at 122°. 5, and in less than 30 parts of boiling water. 
Lead chloride boils between 8G0° and 1000°. Its observed gas 
density is 140; theory requiring 139 for PbCl n . 

Lead chloride and oxide combine to form several oxy- 
chlorides. The hydroxychloride, PbOHCl, obtained by treat- 
ing lead chloride with lime water, is used as a white pig- 
ment. 



388 THE FOURTH GROUP. 

On passing chlorine into a mixture of concentrated hydro- 
chloric acid and lead chloride a red solution is obtained, from 
which lead dioxide separates on addition of water. Lead di- 
oxide forms with concentrated hydrochloric acid a similar 
red solution. These reactions indicate the formation of lead 
tetrachloride, PbCl 4 . 

Lead Iodide. Pbl 2) separates as a yellow precipitate when a soluble 
iodide is added to a solution of a lead salt. From hot solutions it crys- 
tallizes on cooling in golden leaflets. 

Lead Suboxide, Pb 2 0, is obtained as a black powder by heating lead 



It is decomposed at a dull red heat into lead and lead oxide. Acids 
effect the same decomposition, the oxide reacting with the acids to form 

salts. 



Lead Oxide, PbO or Pb = 0, is formed when lead is heated 
in air, the metal being first coated with the suboxide, which 
is soon converted into the oxide. If the lead is constantly 
stirred and the temperature kept below the melting point of 
the oxide a gray ash is formed, consisting of oxide mixed with 
finely divided metallic lead, which is completely converted 
into oxide by continued heating. The oxide thus prepared, 
and known as massicot, has a bright yellow color. If lead is 
heated in air to a temperature at which the oxide fuses, the 
latter flows from the surface of the metal as fast as formed. 
Tbe molten oxide solidifies to a yellowish or reddish mass of 
soft scales which is known as litharge. It is obtained in large 
quantities in cupellation of silver lead alloys. It is used in 
the manufacture of lead salts, in the preparation of drying 
oils, and for a cheap glaze on earthenware. 

Lead oxide absorbs carbon dioxide (11 per cent has been 
found) and moisture from the air. It is soluble in a solu- 



LEAD. 389 

tion of potassium hydroxide, and from the hot concentrated 
solution yellow rhombic crystals of the oxide separate on cool- 
ing. Red tetragonal crystals are obtained when a molten 
mixture of lead oxide and potassium hydroxide is slowly 
cooled. 

Lead Oxyhydroxide, HO-Pb-0-Pb-OH, is obtained as a 
white precipitate when an alkali hydroxide is added to a so- 
lution of lead acetate. If ammonia is used, the compound 
3PbO.H 2 is said to be formed. Lead oxyhydroxide is soluble 
in alkali hydroxides, and to a slight extent in water. The 
aqueous solution absorbs carbon dioxide from the air, and 
becomes turbid. 

Red Lead or Minium. — When lead oxide is carefully 
heated in air it takes up from 1.5 to 2.5 per cent of oxygen, 
and the color changes to a bright red. The product may be 
regarded as a mixture of compounds of lead oxide and diox- 
ide. Red lead is largely used as a pigment, and in the manu- 
facture of glass. It is converted by ignition into oxide. 

Lead Dioxide, Pb0 2 , is obtained as a brown powder by di- 
gesting red lead with dilute nitric acid. It is manufactured 
on a large scale by cheaper methods. 

Exp. 231. — Add red lead gradually, and with constant stirring, to 
nitric acid diluted with its bulk of water. Collect the oxide on a filter, 
wash and dry it. Lead nitrate may he obtained in crystals by evapo- 
rating the acid filtrate. 

Lead dioxide loses half its oxygen on heating, and acts as a 
powerful oxidizing agent towards combustible matter. It has 
feeble basic peoperties, and forms the salt PbO(0,H 3 0.,) 2 , 
which is soluble in glacial acetic acid, and is decomposed by 
water into Pb0 2 and acetic acid. Lead dioxide dissolves in 
very concentrated potassium hydroxide, and the solution 
yields on evaporation in vacuum hydrous crystals of pot as- 



390 THE FOURTH GROUP. 

sium plumbate, = Pb<Q~-jr- + 3H 2 0. This compound is 
analogous to carbonates. 

Lead Sulphide, PbS, occurs as galena, crystallized in cubes 
or modifications of cubes of a bluish-gray color. Hydrogen 
sulphide precipitates from solutions of lead the sulphide as an 
amorphous black mass. In presence of hydrochloric acid a 
reddish precipitate of lead chlorosulphide is sometimes formed, 
which is converted into lead sulphide by an excess of hydro- 
gen sulphide. Lead sulphide dissolves in concentrated hydro- 
chloric acid, hydrogen sulphide being evolved. Nitric acid 
converts it into lead nitrate and sulphate. 

Lead Acetate, Pb(C 2 H 3 2 ) 2 or cHC00 >Pb,— Tllis salt is 
manufactured in large quantities by dissolving lead oxide in 
acetic acid. It is soluble in 1.5 parts of water at ordinary 
temperature and 0.5 part above 100°. The hot solution de- 
posits on cooling crystals containing 3 molecules of water. 
Solutions of lead acetate dissolve lead oxide, with formation of 
basic acetates. Solutions of these basic compounds rapidly 
absorb carbon dioxide from the air, and become turbid from 
separation of lead carbonate. Lead acetate has a slightly 
sweetish taste, and is commonly known as sugar of lead. It 
is used extensively in the arts, in the preparation of lead chro- 
mate, and other pigments. 

Lead Nitrate, Pb(N0 3 ) 2 or N0 2 -0 >Pb— This salt > made b ^ 
dissolving lead oxide in dilute nitric acid, is very soluble in 
water, but insoluble in strong nitric acid. 

Lead Sulphate, PbS0 4 or S0 2 <Q>Pb.— This salt separates 

as a heavy white powder when sulphuric acid or a soluble sul- 
phate is added to a solution of a lead salt. It is soluble in 






LEAD. 391 

22,816 parts of water, and in 36,500 parts of dilute sulphuric 
acid. It is more soluble in the concentrated acid, but sepa- 
rates on diluting the acid solution with water. Lead sulphate 
is somewhat soluble in dilute nitric acid, and readily soluble 
in ammonium acetate. 

Lead Carbonate, PbC0 3 , occurs as the mineral cerusite in 
crystals isomorphous with arragonite. It is obtained arti- 
ficially by pouring a solution of lead nitrate into ammonium 
carbonate. The precipitate produced by sodium carbonate in 
solutions of lead salts is a basic carbonate of variable compo- 
sition. 

White Lead. — This substance is a mixture of basic car- 
bonates which may be viewed as lead carbonate with varying 
proportions of lead hydroxide. It yields on ignition from 85 
to 86.3 per cent of lead oxide. White lead forms with a dry- 
ing oil an opaque white mixture, which is remarkable for its 
covering power, and is the best basis for paints. White lead 
is largely adulterated with barytes (barium sulphate). 

There are several processes employed in the manufacture of 
white lead, the Dutch process being the oldest. In this pro- 
cess sheets of lead in the form of a spiral are placed in earthen 
pots containing dilute acetic acid or vinegar. A number of 
pots are piled together and covered with horse manure or 
spent tan-bark. The vapor of acetic acid and the ox\ T gen of 
the air attack the lead, a basic acetate being formed. The 
fermentation of the bark or manure supplies carbon dioxide, 
Which converts the basic acetate into basic carbonate. The 
acetic acid set free converts a further portion of lead into ace- 
tate, which in turn is decomposed, and hence but little acid 
is required in the process. After some weeks the white lead 
is detached from the metallic lead remaining, and ground 
while moist to a paste, washed to remove acetic acid, and then 
dried. 



392 



THE FOURTH GROUP. 



In the French process a solution of basic lead acetate is 
prepared by digesting litharge with a solution of lead acetate. 
Carbon dioxide is then passed into the solution. The basic 
acetate is decomposed, with separation of white lead, which is 
allowed to settle. The supernatant liquor, containing lead 
acetate and free acetic acid, is drawn off and treated with 
litharge for a new solution of basic acetate. The white lead 
is washed by decantation and dried. 



The student may easily prepare the lead salts of hydro- 
chloric, nitric, and sulphuric acids, and may verify by experi- 
ments some of the statements made about these salts, and also 
about the oxides and the sulphide of lead. He should formu- 
late the reactions involved in the experiments. 



Summary of the Fourth Group. 

Atomic Weight. Density. Fusing Point. 

Carbon, 12 3.5 Infusible. 

Silicon, ..... 28 2.5 White heat. 

Titanium, .... 48 — 

Zirconium, .... 90.7 4.1 Infusible. 

Cerium, 141 6.6 Below 1000°. 

Thorium, 232 11 Infusible. 

Germanium,. ... 72 5.5 900°. 

Tin, 118 7.3 230°. 

Lead, 207 11.35 334°. 

The densities of the elements of this group increase with 
the atomic weights from carbon and silicon to thorium, and 
from germanium to lead. Most of them are infusible, or 
fusible only at very high temperatures. Tin and lead, how- 
ever, are readily fusible. While the members of the group 
present few analogies in physical properties, they exhibit to a 



SUMMAKT. 393 

striking degree in their compounds a similarity in chemical 
characteristics. With the halogens all the members of the 
group except cerium form tetrahalides, and all form dioxides. 
Carbon is acidic in carbonates, but with hydrogen it forms 
basic radicals. Silicon is acidic, while titanium with the next 
higher atomic weight is both acidic and basic. Zirconium 
has more pronounced basic characters, and cerium and tho- 
rium exhibit only basic properties. 

Tin forms acid hydroxides, and the salts which the metal 
forms with oxygen acids have little stability. Lead, on the 
contrary, is strongly basic, and forms well-defined salts con- 
taining bivalent lead. 

Cerium forms by preference compounds analogous to those 
of the third group, as for example, CeCl 3 and Ce 2 (S0 4 ) 3 . 

The student should make a table of the formulas of the 
compounds of the group which have been described, placing 
analogous compounds in the same column. Such a table will 
exhibit the chemical analogies existing between the different 
members of the group. 



THE EIGHTH GKOUP. 



The members of this group are iron, cobalt, and nickel, 
and the platinum metals, which are ruthenium, rhodium, 
palladium, platinum, iridium, and osmium. The platinum 
metals are so called because of their similarity to platinum, 
with which they are associated in nature, all occurring in the 
metallic state. 



Iron (Ferrum), Fe. 

Atomic Weight, 56. Density, 7.84. 

Iron is the most useful of metals. No other metal pos- 
sesses such a variety of valuable properties, or can be used 
instead of it for cutting-tools. Iron is widely distributed, 
occurring in all soils and most rocks, and is a constituent of 
many minerals. It is found in small quantities in animals 
and plants, and in many natural waters. The minerals limo- 
nite, 2Fe 2 0~ 3 -f- 3H 2 0, and oxides of iron with less water, 
also hematite, Fe 2 3 , magnetite, Fe 3 4 , and siderite, FeC0 3 , 
constitute the ores of iron. Some of the ores are nearly pure, 
but most of them are mixed with earthy matter, limestone, 
quartz, and compounds of manganese. Sulphur and phos- 
phorus are usually found in small quantities in the ores, 
and are troublesome impurities, owing to difficulty in separat- 
ing them from the iron. 

The different kinds of iron used in the arts are comprised 
under three heads, viz.: Wrought Iron, Steel, and Cast Iron. 
Their differences are exhibited in part by the following table: 



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396 THE EIGHTH GKOUP. 

Manufacture of Cast Iron. — Most of the iron of com- 
merce is obtained by the reduction of ores to cast or pig iron 
in blast furnaces. These furnaces vary greatly in size and 
shape, but are essentially shafts narrowed at the top and 
bottom, and having a height of 30 to 90 feet. Their capacity 
varies from 4000 to 40,000 cubic feet. The furnace is 
charged at the top with ore and fuel, the latter being either 
hard coal, coke, or charcoal. Limestone is usually added as a 
flux to form a fusible slag with the impurities of the ore. A 
blast of air is blown through four or more openings called 
tuyeres, near the bottom of the furnace. The oxygen of the 
air and the incandescent carbon form carbon monoxide, which 
passes upwards, and near the top of the furnace reduces the 
lumps of the ore to porous masses of unfused iron. The 
carbon monoxide takes the oxygen from the oxide of iron, 
carbon dioxide being formed. As the iron sinks in the fur- 
nace it is exposed to higher temperatures, and very compli- 
cated reactions occur between the iron and the oxides of 
carbon, and perhaps other gases, and the fuel. At this stage 
of the process the iron takes up carbon and silicon, and is 
converted into cast iron, which fuses and sinks to the bottom 
of the furnace. The impurities of the ore unite with the fltix 
to form a liquid slag, which floats on top of the liquid iron, 
and flows from an opening in the side of the furnace. The 
iron is drawn off from time to time, and cast into pigs. 

Varieties of Cast Iron. — The kind of iron made in a blast 
furnace depends upon the quality of the ore, and partly upon 
the amount of fuel and blast used. Gray irons contain carbon, 
partly in combination with iron, and partly as graphite me- 
chanically mixed. In white iron nearly all the carbon is in 
combination. Intermediate between these two kinds are the 
mottled irons. Gray iron is used for ordinary castings. 
Spiegeleisen is a variety of white iron which is very hard and 
brittle, containing considerable manganese, and more carbon 



IROK. 397 

than other varieties of iron. Certain qualities of gray iron 
when suddenly cooled from the molten state are converted 
into white iron, and are said to chill. Chilled iron is 
harder and resists wear better, but is more brittle, than ordi- 
nary gray iron. Cast-iron car-wheels are made of a good 
quality of iron, cast in a mould with a band of iron to chill 
the rim of the wheel to a depth of about an inch ; the 
remainder of the wheel, cooling slowly in the sand, is tough 
gray iron. 

Manufacture of Wrought Iron. — The ores of iron are re- 
duced by heating in a suitable fire or furnace to a sponge-like 
mass of wrought iron. This method, which is costly in fuel, 
and does not admit of production on a large scale, was proba- 
bly the earliest method, and is still employed by natives of 
India and Central Africa. Wrought iron is commonly made 
by burning out the carbon, silicon, sulphur, and phosphorus 
of cast iron by the puddling process. In this process pig iron 
is melted by an oxidizing flame on the hearth of a furnace, 
and stirred or puddled in order to expose fresh surfaces of 
the molten metal to the action of oxygen, some iron oxide 
being added to facilitate the oxidation of the impurities. 
The iron gradually becomes less fusible, and at last changes 
to a pasty mass. It is then gathered into lumps, which are 
removed from the furnace, and squeezed and hammered to 
get rid of slag, and finally rolled into bars. The iron is 
made more homogeneous and tougher by heating bundles of 
bars to a welding heat and rolling again. 

Manufacture of Steel. — It is not found practicable to make 
steel directly from pig iron by burning out only part of the 
carbon. In the manufacture of steel by the cementation 
process bars of wrought iron are heated for some days in con- 
tact with charcoal powder. The iron slowly takes up car- 
bon and changes to steel. The metal is not uniform in com- 
position or structure, and to overcome these imperfections it is 



398 THE EIGHTH GROUP. 

melted in crucibles and cast into ingots. The product at this 
stage is coarsely crystalline, and contains cavities. The ingots 
are heated and rolled, and the steel is thus rendered fine- 
grained and tougher. The finest qualities of tool steel are 
made by the cementation process. 

The Bessemer method of making steel is much less costly 
than the one just described. In this process molten cast iron 
is run into a large vessel called a converter, which is lined 
with an infusible mixture of clay and quartz, and a powerful 
blast of air is blown into the molten metal through a number 
of small holes in the bottom of the converter. The silicon, 
carbon, and other impurities burn out rapidly, leaving nearly 
pure iron. Spiegeleisen is then added in order to introduce 
the requisite amount of carbon, and the metal is poured 
from the converter into moulds. The ingots of steel thus 
produced are afterwards rolled while hot. 

The Bessemer process is well adapted for making low steels, 
i.e., steels with little carbon, which have largely replaced 
wrought iron for railroad rails, boiler-plates, and other pur- 
poses. Formerly all varieties of Bessemer metal were called 
steel, but at the present time if the carbon content is so low 
that the metal will not harden it is called ingot iron. 

In the so-called " basic process" the converter has a lining 
chiefly of lime. In the older processes a portion of the phos- 
phorus in the cast iron remains in the steel, whereas in the 
basic process the phosphorus is removed by the lime. 

There are other methods of making steel, such as melting 
together wrought and cast iron in crucibles, and by smelting 
ores on a hearth and then adding spiegeleisen. 

Preparation of Pure Iron. — Pure iron is little known. It 
has been obtained by various methods, one of which is the 
following. Pure iron oxide is reduced by hydrogen in a 
platinum crucible, and the finely divided iron thus obtained 
is then fused in a lime crucible by means of an oxyhydrogen 
blow-pipe supplied with pure gases. 



iron. 399 

Properties of Iron and Steel. — These have been given in part 
in the table on page 395. Pure iron is almost silver white, and 
possesses greater tensile strength than any other metal, except- 
ing nickel and cobalt. It is softer, tougher, and has a higher 
melting point than ordinary wrought iron, which it resembles 
in other properties. 

Wrought iron possesses the valuable property of softening at 
a red heat, when it can be forged or rolled into any desired 
shape. At higher temperatures it welds ; that is, pieces with 
clean surfaces adhere and form a homogeneous mass when 
hammered together. Steel does not weld readily, but steel 
and wrought iron are easily welded together. 

The tensile strength of puddled iron varies from 20 to 30 
tons per square inch ; ingot Bessemer iron from 30 to 40 tons ; 
steel is stronger, and steel wire has been made with a break- 
ing strain of 92 tons per square inch of cross-section. 

Sulphur, other than in minute traces, renders iron and steel 
brittle when hot (red short), and phosphorus brittle when 
cold (cold short). These effects are modified to some extent 
by manganese, carbon, and silicon. A steel with 0.5 per cent 
of carbon and 0.1 per cent of sulphur is liable to be red short, 
and 0.1 per cent of phosphorus makes it cold short. Even 
0.03 per cent of phosphorus injures tool steel. 

Steel which has been heated to faint redness, and cooled by 
plunging into water, is very hard, and too' brittle for use. To 
make it tougher the temper is drawn by heating, the progress 
of the change being judged from the tints on the polished 
surface due to oxidation. At about 230° the steel becomes a 
light straw yellow, and is still very hard ; at 243° the color is 
full yellow, and the temper is suitable for razors ; at 255° a 
brownish yellow appears, and the temper is adapted to cold 
chisels ; at 265° purple first appears, and the temper is suit- 
able for axes; at 288° bright blue is scon, which is the color 
for watch-springs ; at higher temperatures up to 31 l>° softer 
tempers result. Steel and iron are annealed by slow cooling. 



400 THE EIGHTH GKOUP. 

Iron does not change in dry air, nor in water free from air, 
but when exposed to common air or water it rusts, and the 
rusting is hastened by the presence of acids, even carbonic. 
Lime water and other alkaline solutions prevent the rusting of 
iron. A highly polished surface resists corrosion longer than 
a rough surface. Any covering, such as paint, oil, or graphite, 
prevents rusting. A very effective protection is obtained by 
heating iron and steel articles in air, carbon dioxide, or steam, 
until a coating of magnetic oxide is formed. The so-called 
"galvanized iron" is iron coated with zinc by immersion in a 
bath of molten zinc. The coating of zinc prevents the iron 
from rusting. 

Iron when hot burns readily in air and oxygen, and also in 
sulphur vapor. It unites with the halogens at ordinary tem- 
perature. White cast iron is less readily soluble in acids than 
gray cast iron or wrought iron. 

Iron forms two well-defined classes of compounds, viz., 
ferrous, in which Fe is apparently bivalent, as in FeCl 2 ; and 
ferric, in which Fe is assumed to be trivalent, as in FeCl 3 . 
Iron is also regarded by chemists as tetravalent, and the con- 
stitution of the two series of salts has been expressed by such 
formulas as 

~ 1 1 >Fe=Fe<™' and Cl-^Fe-Fe^-Cl. 

ci 01, C1 / \ 0l 

In this book the simpler formulas are used. 



Ferrous Compounds. 

CI 
Ferrous Chloride, FeCl 2 or Fe < «,, is obtained in colorless 

scales when hydrochloric acid gas is passed over red-hot iron 
wire, and also by reducing ferric chloride in a current of hy- 
drogen. It is very deliquescent, fuses at a red heat, and vola- 



FERROUS COMPOUNDS. 401 

tilizes only at a high temperature. Its gas density has been 
found to be 92.5 at a yellow heat; FeCl 2 requires 63.2 and 
Fe 2 Cl 4 126.4 : hence it appears that Fe 2 Cl 4 exists, decomposing, 
however, at higher temperatures into FeCl 2 . The hydrate, 
FeCl 2 -f 4H 2 0, is obtained in bluish crystals when a hot con- 
centrated solution of iron in hydrochloric acid is cooled out of 
contact with air. 

Ferrous Iodide, Fel 2 or Fe< T , is formed by the direct union 

of its elements. A solution of the salt is easily prepared by 
digesting iron filings with iodine and water. The colorless or 
bluish-green solution absorbs oxygen from the air, and decom- 
poses with separation of iodine and formation of ferric hy- 
droxide. Ferrous iodide is used in medicine. 

Ferrous Oxide, FeO or Fe=0. — This oxide is obtained by 
heating ferric oxide to 500° in a current of hydrogen. It is a 
black powder, which rapidly oxidizes in air. 

OH 

Ferrous Hydroxide, Fe(0H) 2 or Fe<Qjx. — An alkali hy- 
droxide produces in a solution of a pure ferrous salt a white 
flocculent precipitate of ferrous hydroxide. If the precipitate 
is washed out of contact with air and dried in hydrogen, the 
hydroxide remains in the form of a white powder. Ferrous 
hydroxide rapidly absorbs oxygen. It dissolves in about 
150,000 parts of water, forming an alkaline solution having 
the taste of ferrous salts. 

Ferrous Nitrate, Fe(N0 3 ) 2 or ^0 -0 >Fe,— This salt is ob ~ 

tained by dissolving ferrous sulphide in cold dilute nitric acid 
(density 1.12). The solution deposits on evaporation in va- 
cuum at low temperatures crystals of Fe(NO s ) a + 6H a O. A 
solution of ferrous nitrate, free from acid, may be boiled with- 
out decomposition, but in presence of free nitric acid it is con- 



402 THE EIGHTH GEOUP. 

verted into ferric nitrate, and nitric oxide is given off. Iron 
dissolves in cold dilute nitric acid without evolution of hydro- 
gen, and the solution contains ferrous nitrate and ammonium 
nitrate. Such a solution is used by dyers. 

Ferrous Sulphate, FeS0 4 or S0 2 < Q >Fe.— A solution of this 

salt is obtained by dissolving iron in dilute sulphuric acid. 
On cooling a concentrated solution, and on evaporating a 
dilute solution, bluish-green monoclinic crystals of FeS0 4 
+ 7H 2 separate, which are known as green vitriol. Ferrous 
sulphate is very soluble, 100 parts of FeS0 4 + ~H 2 dissolving 
in 164 parts of water at 10° and 36 parts at 100°. Ferrous 
sulphate may be crystallized with five and with four mole- 
cules of water. The monohydrate, FeS0 4 -f H 2 0, is obtained 
as a white powder by heating green vitriol to 110° out of con- 
tact with air ; at a little higher temperature the salt becomes 
anhydrous. Green vitriol, when pure and dry, does not 
absorb oxygen from the air unless the latter is very moist. 
The commercial salt commonly contains ferric sulphate, which 
changes the delicate bluish-green color of the pure salt to 
green. Crystals containing ferric sulphate become moist on 
exposure and slowly absorb oxygen, with formation of a red- 
dish-yellow coating of basic ferric sulphate. Solutions of 
ferrous sulphate undergo the same change, and deposit a 
yellow basic ferric sulphate. 

Green vitriol is manufactured in large quantities from iron 
pyrites, FeS 2 , and it is a by-product in iron-works where iron 
is pickled or cleansed with sulphuric acid. It is used in dye- 
ing, in purifying coal gas, and as a disinfectant. 

Ferrous Ammonium Sulphate, FeS0 4 .(NH 4 ) 2 S0 4 + 6H 2 0.— 

This salt separates in light bluish-green crystals on cooling 
a concentrated solution of green vitriol and ammonium sul- 
phate, mixed in the requisite proportions. It is much more 
permanent than green vitriol, and on this account, and the 



FERROUS COMPOUNDS. 403 

ease with which it is obtained pure, is useful in standardizing 
solutions required in the volumetric estimation of iron. It 
contains one seventh of its weight of iron. 

Ferrous Carbonate, FeC0 3 , occurs abundantly as an ore 
known as spathic iron. It is obtained artificially by heat- 
ing to 150° for 12 to 36 hours a solution of ferrous sul- 
phate to which sodium carbonate has been added. Thus 
prepared it is a white crystalline powder, which is not 
readily acted upon by acids, and is permanent in dry air. 
The hydrated ferrous carbonate separates on addition of 
an alkali carbonate to a ferrous solution as a white floc- 
culent precipitate, which absorbs oxygen rapidly, gives off 
carbon dioxide, and finally changes to ferric hydroxide. If 
the precipitate is thoroughly washed with deaerated water, 
and dried in carbon dioxide, a greenish tasteless powder is ob- 
tained, which is tolerably permanent. 

Ferrous carbonate and metallic iron dissolve to a slight 
extent in water containing carbon dioxide. Many natural 
waters contain iron held in solution by carbonic acid, and are 
known as chalybeate waters. The solution has a metallic 
flavor, and decomposes on exposure to air, with separation of 
ferric hydroxide. Solutions of ferrous carbonate have been 
recently used to purify sewage water. 

Ferrous Sulphide, FeS.— This compound is made by drop- 
ping sulphur upon red-hot scrap iron. The iron and sulphur 
melt together, and on allowing the product to cool a brittle 
gray mass of ferrous sulphide is obtained. It is, however, 
liable to contain particles of uncombined iron. If sulphur is 
added to the hot sulphide, until the latter on cooling is yellow, 
owing to the formation of a higher sulphide, and the mass is 
then fused, a pure product is obtrned. Ferrous sulphide is 
much used for the preparation of hydrogen sulphide. An 
amorphous hydrous ferrous sulphide is formed when ammo- 
nium sulphide is added to a ferrous solution. If a ferric salt 



404 THE EIGHTH GEOUP. 

is used, the precipitate contains free sulphur. The reaction 
with ferric chloride is 

2FeCl 3 + 3(NH 4 ) 2 S = 2FeS + S + 6NH 4 C1. 

Hydrous ferrous sulphide absorbs oxygen rapidly on exposure 
to the air, and is converted into ferrous sulphate. 

Iron Disulphide, FeS 2 , is not a ferrous compound, but is 
conveniently noticed in connection with ferrous sulphide. It 
occurs as the mineral pyrite or pyrites in fine yellow crystals 
belonging to the cubic system. On heating FeS 2 half of the 
sulphur is given off, leaving ferrous sulphide. Pyrites is 
largely used for its sulphur in the manufacture of sulphuric 
acid. 

Exp. 232. — a. Place in a flask 10 cc. of concentrated sulphuric acid, 
30 cc. of water, and about 10 grams of fine iron wire. "Warm gently, 
and after the reaction has nearly ceased filter the hot solution into a 
porcelain dish. On cooling, crystals (of what ?) will separate. Pour off 
the mother-liquor, and rinse the crystals once with a small quantity of 
water, b. Dry a portion of the crystals with filter paper, c. Leave 
another portion moist, d. Dissolve the remainder in water. Set all 
three preparations aside for a few days, and note any changes. Repre- 
sent by an equation the reaction between the iron and acid, and state 
observations which support the view expressed by the equation. 

Exp. 233. — Treat iron wire in a test-tube with dilute sulphuric acid. 
When considerable has dissolved pour the solution while boiling hot 
into a test-tube containing boiling dilute ammonia. The latter should 
be in excess. Formulate the reaction, and note color of precipitate and 
any changes it may undergo on standing. 

Exp. 234. — Take 20 grams of green vitriol — the commercial article 
will answer — and the weight of common ammonium sulphate found by 
calculation to be required to form ferrous ammonium sulphate. Dis- 
solve the two salts together in 50 cc. of boiling water, and filter the hot 
solution into a wide-mouthed bottle. Label and set aside. Crystals, 
often of considerable size, will separate after a day. Test some of the 
crystals for iron and for ammonia. Write equations of the reactions by 
which iron and ammonia are found. 

Exp. 235.— Add ammonium sulphide to a solution of ferrous sulphate, 



FEERIC COMPOUNDS. 405 

and wash the precipitate obtained on a filter several times with water. 
Leave a portion of the moist ferrous sulphide in a porcelain dish until 
next practice, when any changes may be noted. Treat another portion 
of the precipitate with dilute hydrochloric acid. What is the result ? 



Ferric Compounds. 

/CI 

Ferric Chloride, Fed, or Fe(-Cl. — Hot iron burns in 

\C1 
chlorine with the formation of this chloride. Chlorine con- 
verts ferrous chloride into the ferric salt. When a solution of 
ferric chloride is evaporated to dryness the dry residue con- 
sists of ferric chloride mixed with oxychloride, and on gentle 
ignition the former sublimes and condenses in the cooler parts 
of the apparatus. Ferric chloride forms grayish-black leaf- 
lets, very soluble in water and alcohol, and which deliquesce 
in air to a dark-red liquid. 

The observed gas density of ferric chloride at 448° is 151.6 ; 
theory requires 162.2 for Fe 2 Cl 6 . At higher temperatures 
ferric chloride is partially dissociated into ferrous chloride 
and chlorine. It is possible that molecules of FeCl 3 and of 
Fe 2 Cl 6 exist in the gaseous state. 

A solution of ferric chloride is easily made by dissolving- 
iron in hydrochloric acid, and then passing chlorine gas into 
the solution to convert the ferrous into ferric chloride ; or by 
dissolving ferric oxide or hydroxide in hydrochloric acid. A 
concentrated solution of ferric chloride is brownish red. the 
color changing to a faint yellow on diluting the solution. The 
hydrate FeCl 3 -(- 6H 2 is obtained as a solid crystalline mass 
when 100 parts of anhydrous ferric chloride are mixed with 
63.5 parts of water. 

Ferric Oxide, FeO, or 0=Fe-0-Fe = 0, is also called sesqui- 

oxide of iron and red oxide of iron. It occurs as the mineral 



406 THE EIGHTH GROUP. 

hematite, which is one of the principal ores of iron. It is ob- 
tained artificially by igniting any of the ferric hydroxides or 
salts of iron with volatile acids. Ferric oxide of whatever 
source gives a red powder when pulverized. It is used in 
paint, and in the form of a powder called rouge is specially 
prepared for polishing glass. That used for polishing metals 
is called crocus. Strongly ignited ferric oxide dissolves slowly 
in hydrochloric and sulphuric acids. 

Ferrous-Ferric Oxide, Magnetic Oxide of Iron, Fe 3 4 or 
= Fe-0-Fe-0-Fe = 0. — This oxide occurs as magnetite, crys- 
tallized in octahedrons. It is widely distributed, and is the 
richest ore of iron, containing when pure 72 per cent of iron. 
It is easily distinguished by being strongly attracted by the 
magnet, and by yielding a black powder. 

Some native magnetites, known as lodestones, possess the 
property of attracting iron and rendering it magnetic. The 
magnetite ores are not commonly lodestones, but often will 
lift small pieces of fine iron wire. 

The magnetic scale formed when iron is heated in air is a 
varying mixture of ferrous and ferric oxide, the ferrous oxide 
predominating in the inner portion and the ferric oxide on 
the outside. Eeference has already been made to the forma- 
tion of magnetic oxide when hot iron and steel are exposed to 
steam or carbon dioxide. On the other hand, all the iron ox- 
ides are reduced to the metal by hydrogen and carbon monox- 
ide at red heat. 

When a solution of magnetic oxide in hydrochloric acid or 
a solution of ferrous and ferric salts in the right proportion is 
treated with ammonia, a black magnetic hydroxide is precipi- 
tated. 

Ferric Hydroxides. —Ammonia produces in cold rather 
dilute solutions of ferric salts a bulky brownish-red precipitate 
of ferric hydroxide, Fe(OH) 3 . This readily loses water, and 



FERRIC COMPOUNDS. 40? 

in boiling water becomes less bulky and darker colored. The 
dehydration may thus be carried so far that only 2 per cent of 
water will remain with the precipitate. Freshly precipitated 
ferric hydroxide dissolves readily in acid, but when more or 
less dehydrated dissolves slowly. Iron rust has the compo- 
sition Fe 2 3 + 2Fe(OH) 3 . 

Soluble Ferric Hydroxide. — Freshly precipitated ferric hy- 
droxide dissolves in a solution of ferric chloride, forming a 
deep-red liquid. A similar solution is obtained when ammo- 
nium carbonate is added to ferric chloride until the precipi- 
tate formed no longer dissolves on shaking. If such a solution 
is dialyzed a red solution remains, which is nearly free from 
chlorides, and from which after some weeks gelatinous ferric 
hydroxide separates. 

N0 2 -(k 
Ferric Nitrate, Fe(N0 3 ) 3 or NO -0~Fe.— When iron dis- 

NO-CK 
solves in nitric acid, density 1.034, ferrous nitrate is formed, 
together with ammonium nitrate, as stated under ferrous 
nitrate. If, however, acid having a density of 1.115 is taken, 
only ferric nitrate is formed, and no ammonium salt. Crys- 
talline hydrates of ferric nitrate have been obtained from con- 
centrated solutions. When a solution of ferric nitrate is 
evaporated in a water-bath, nitric acid is given off, and a 
residue remains from which water dissolves little, and which 
is slowly taken up by hot nitric acid. Ferric nitrate is used 
in dyeing. 

S0 =<0)Fe 
Ferric Sulphate, Fe o (S0 4 ) 3 or S0,<q .—This salt is 

easily prepared by adding sulphuric and nitric acids to a solu- 
tion of ferrous sulphate : 
6FeS0 4 + 3II 2 S0 4 + 2IIN0 3 = 3Fe 9 (S0 4 ) 8 + 2X0 + 411 .0. 






408 THE EIGHTH GEOUP. 

When a solution containing an excess of sulphuric acid is 
evaporated, the salt separates as a white powder, which dis- 
solves with difficulty in water. The presence of a small quan- 
tity of ferrous sulphate facilitates the solution. A number of 
complex basic ferric sulphates have been described. 

Ferric Ammonium Sulphate, Fe 2 (S0 4 ) 3 .(NH 4 ) 2 S0 4 + 24H 2 0.^ 
This double salt is an iron alum. It is made by adding the 
required amount of ammonium sulphate to a solution of ferric 
sulphate containing a little free sulphuric acid. The solution 
deposits on spontaneous evaporation octahedral crystals of 
the alum. It is emplo} r ed in dyeing, when a perfectly neutral 
ferric salt is required. 

Exp 236. — To a solution of green vitriol add some sulphuric acid, 
then drop by drop nitric acid, until on warming red fumes are no 
longer evolved and the green solution has become reddish. Dilute a 
portion of the solution of ferric sulphate with water, and add ammonia 
in excess. Express the reaction by an equation. Boil for some time, 
aud observe that the precipitate becomes denser and darker colored. 
Wash the precipitate thoroughly on a filter with hot water, dry at 100°, 
and, after removing from the filter, ignite in a porcelain crucible. The 
ferric oxide thus prepared will yield a red powder when rubbed in a 
mortar. 

Ferric Acid, H 2 Fe0 4 , is not known in the free state. When fine iron 
filings are fused with two parts of potassium nitrate, the mass after 
cooling yields with water a red solution, which contains potassium 
ferrate. The same salt is also formed when chlorine is passed through 
a solution of potassium hydroxide in which ferric hydroxide is sus- 
pended. Potassium ferrate is unstable, decomposing readily in aqueous 
solution, with separation of ferric hydroxide and evolution of ox3 r gen. 
Alkaline solutions are more permanent. Barium ferrate, BaFe0 4 + 
H 2 0, separates as a dark-red powder when barium chloride is added to 
a solution of the potassium salt. 



Iron Cyanides. 



Silver cyanide dissolves in a solution of potassium cyanide, 
forming the double salt AgCN.KCN. Nitric acid decom- 



IKON CYANIDES. 409 

poses this compound, precipitating silver cyanide and setting 
free hydrocyanic acid, thus : 

AgCN.KON + HN0 3 = AgCN + HON + KN0 3 . 

Hydrogen sulphide separates the silver from the cyanogen: 

2(AgCN.KCN) + H 2 S == Ag 2 S + 2HCN + 2KCN. 

Potassium cyanide produces in solutions of zinc salts a white 
precipitate of zinc cyanide, which forms with potassium cy- 
anide the soluble salt 2KCN.Zn(CN) 2 . This salt is decom- 
posed by nitric acid, with separation of all of the cyanogen as 
hydrocyanic acid : 

2KON.Zn(CN) a + 4HTO 3 = 4HCST + 2KNO, + Zn(N0 3 ) 2 . 

Ammonium sulphide precipitates the zinc as sulphide : 

2KON.Zn(ON) 9 + (NH 4 ) a S = ZnS + 2KCJST + 2NH 4 CN. 

A number of double cyanides are decomposed with separa- 
tion of cyanogen from the metals by reactions analogous to 
those above given. There are, however, compounds contain- 
ing iron and cyanogen, known as ferro- and ferricyanides, 
which exhibit a different deportment. The iron .and cyano- 
gen in them are not separated from each other by reagents 
which precipitate iron from its ordinary salts, such as ferrous 
sulphate or ferric chloride, or which separate cyanogen from 
potassium, zinc, or silver cyanide. 

Potassium cyanide produces in a solution of ferrous sul- 
phate a yellowish-red precipitate, soluble in an excess of po- 
tassium cyanide, with formation of potassium ferrocyanide, 
K 4 (0 3 N 3 ) 2 Fe. From solutions of this salt acids fail to separate 
ferrous cyanide, or at common temperature to set free hydro- 
cyanic acid. Moreover, neither ammonium sulphide nor al- 
kalies separate the iron from potassium t'errocvanide. The 
potassium in K 4 (C 3 N 3 ).,Fe can be replaced by hydrogen, and 



410 THE EIGHTH GROUP. 

by metals such as lead, calcium, and zinc, "without, however, 
separating the iron from the cyanogen. These reactions show 
that iron and cyanogen are united in a compound radical, 
which can be transferred from one compound to another with- 
out losing its identity. 

Assuming that the precipitate which potassium cyanide pro- 
duces in a solution of ferrous sulphate is ferrous cyanide, a 
compound not yet obtained pure, we may represent the for- 
mation of potassium ferrocyanide thus : 

Ee(CN) 2 + 4KCT = K 4 (C 3 N 3 ) 2 Fe. 

Tricyanogen chloride, C 3 N 3 C1 3 , is the acid chloride of cy- 
anuric acid, C 3 N 3 (OH) 3 , both compounds containing the triv- 
alent radical tricyanogen, C 3 N 3 . Oyanuric acid is formed 
when anhydrous potassium ferrocyanide is heated with bro- 
mine, and then heating the product with water to convert the 
C 3 N 3 Br 3 , which is doubtless formed, into the acid (Merz and 
Werth). The formation of cyanuric acid from potassium fer- 
rocyanide is explained by the hypothesis that the ferrocyanide 
contains the radical tricyanogen. There is no evidence of the 
existence of a radical composed of carbon and nitrogen more 
complex than C 8 N 8 . The iron in potassium ferrocyanide is 
assumed to be bivalent, since the compound is formed by the 
combination of ferrous cyanide and potassium cyanide. 

The foregoing facts lead to the formula 



K=(C 3 N 3 ) 
K 2 =(C 3 N 3 ) 



>Fe, 



in which the acid radical is ferrocyanogen, [(C^N^Fe] 1 ^ 
The structure of tricyanogen may be 



■C-BT=C- 




-N=C=]\ X 


ii l 


or 


ii ii 


1 




i 



IROK CYAKIDES. 411 

according as the nitrogen in it is considered to be triyalent or 
pentavalent. 

Ferric chloride and potassium cyanide react to form potas- 
sium ferricyanide, K 3 (C 8 N 8 ) 2 Fe. Assuming the iron in this 
salt to be trivalent, we have the structural formula 

in which the acid radical is ferricyanogen [(C 3 N 3 ) 2 Fe] m . 
Oxidizing agents convert potassium ferrocyanide into ferricy- 
anide, and the latter is changed by reducing agents into ferro- 
cyanide. These reactions are analogous to the changes which 
ferrous and ferric salts undergo with like treatment. 

Potassium Ferrocyanide, K 4 (C 3 N 3 ) 2 Fe + 3H 2 0, is known in 
commerce as yellow prussiate of potash. It is the starting- 
point in the preparation of most 01 the cyanides, and is largely 
used in dyeing. It is manufactured on a large scale by adding 
a mixture of- iron filings or iron oxide and animal substances, 
such as horn, dried blood, or leather clippings, to molten po- 
tassium carbonate. The fused mass after cooling is lixiviated 
with water, and the solution on evaporating yields crystals of 
potassium ferrocyanide. The animal matter contains carbon 
and nitrogen, which react with potassium carbonate to form 
potassium cyanide. The sulphur in the animal matter, and 
impure potassium carbonate unites with the iron, forming 
ferrous sulphide, which reacts with the potassium cyanide 
to form potassium ferrocyanide, when the fused mass is treated 
with water. 

Potassium ferrocyanide forms yellow crystals, which lose 
their water of crystallization at 100°, and change to a white 
powder. The crystals dissolve in two parts of boiling and in 
four parts of cold water. Potassium ferrocyanide fuses at a 
red heat, decomposing into potassium cyanide and a com- 
pound of carbon and iron. 



412 THE EIGHTH GROUP. 

Ferrocyanic Acid, H 4 (C 3 N 3 ) 2 Fe. — When concentrated hydro- 
chloric acid is added to a cold saturated solution of potassium 
ferrocyanide, ferrocyanic acid separates in white scales, which 
may be purified by dissolving in alcohol and then adding ether 
to reprecipitate the acid. Ferrocyanic acid is a strong acid, 
capable of decomposing carbonates, acetates, and oxalates. It 
is an unstable compound, and on exposure to air decomposes 
with formation of Prussian blue. A number of ferrocyanides 
are known. 

Potassium Ferricyanide, K 3 (C 3 N 3 ) 2 Fe, is commonly known 
as red prussiate of potash. It is formed when potassium fer- 
rocyanide in solution is treated with reagents which convert 
ferrous into ferric iron. Thus, when chlorine is passed into a 
solution of potassium ferrocyanide, potassium ferricyanide and 
potassium chloride are formed : 

K 4 (0,N",) 2 Fe + 01 = K 3 (C 3 N 3 ).Je + KCl. 

The two salts are separated by repeated crystallization. 
Potassium ferricyanide forms dark-red crystals. The aqueous 
solution of the salt is yellowish brown when concentrated, and 
yellow when dilute. The solution decomposes slightly on 
long exposure to light, with formation of ferrocyanide and a 
blue precipitate. 

Lead Ferricyanide, Pb 3 ,[(C 3 N 3 ) 2 Fe] 2 -f- 16H 2 0, separates when 
hot concentrated solutions of potassium ferricyanide and lead 
nitrate are mixed. More of the salt is deposited on cooling. 
It is sparingly soluble in cold water. 

Ferricyanic Acid, H 3 (C 3 N 3 )Fe, is prepared by adding the 
required amount of sulphuric acid to a solution of the lead 
salt. The solution filtered from the lead sulphate yields when 
evaporated at common temperature ferricyanic acid in deli- 
cate long brown crystals. 



IEOK CYANIDES. 413 



ill II 



Soluble Prussian Blue, K(C 3 N 3 ) 2 FeFe + Aq.— When solu- 
tions of equal molecules of potassium ferrocyanide and ferric 
chloride are mixed a blue precipitate is formed. Potassium 
ferricyanide and ferrous chloride yield the same result. The 
precipitate in either case is soluble Prussian blue, and, after 
washing with water to remove the potassium chloride formed, 
it dissolves in pure water to a blue solution. From this solu- 
tion it is precipitated by addition of salts. The following equa- 
tions represent the formation of soluble Prussian blue : 

Cc S N/ Fe+FeC1 ' = F i ycS!> Fe + 3Kci ; 

Potassium ferrocyanide Soluble Prussian blue 

nH AT 
i_ c »e + FeCl, = K-(CJ : )> e + 3K0 l: 

Potassium ferricyanide Soluble Prussian blue 

The two formulas are evidently identical, each representing 
ferrous and ferric iron joined to two groups of tricyanogen. 

in ii 

Ferric Ferrocyanide, Fe 7 (CN) 18 or Fe 4 [(C 3 N 3 ) 2 Fe] 3 .— This 

compound, known as Prussian blue, is obtained as a blue pre- 
cipitate when potassium ferrocyanide is added to an excess of 
ferric chloride : 

3K 4 (C 3 N 3 ) 2 Fe + 4FeCl 3 = Fe 4 [(0,N,) 9 Fe], + 12KC1. 

Ferric ferrocyanide is also formed when a ferric salt is 
added to soluble Prussian blue. 

ii in 

Ferrous Ferricyanide, Fe r ,(CN) 12 or Fe 8 [(C 3 N 8 ) a Fe] 9 .— When 

potassium ferricyanide is added to a ferrous salt a dark-blue 
precipitate, known as TurnlmU's blue, is formed : 

3FcCl a + 2K,(C,N,) 3 Fe = Fe 8 [(C 8 N,) 8 Fe], -| 6KC1, 



414 THE EIGHTH GEOUP. 

The same compound is obtained by mixing a solution of a 
ferrous salt with soluble Prussian blue. 

Ferrous Thiocyanate, Fe(NCS) 2 . — Iron dissolves in thiocy- 
anic acid, with formation of ferrous thiocyanate. The salt 
separates in light-green crystals, containing three molecules 
of water, when the solution is evaporated out of contact with 
air. The aqueous solution becomes red on exposure to air. 

Ferric Thiocyanate, Fe(NCS) 3 , is formed and imparts a deep- 
red color to the solution when solutions of a ferric salt and a 
thiocyanate are mixed. Solid ferric thiocyanate is obtained 
by treating a mixture of anhydrous ferric sulphate and potas- 
sium thiocyanate with alcohol, which dissolves out ferric thio- 
cyanate, leaving potassium sulphate. The alcoholic solution 
on evaporation in vacuum deposits dark-red crystals, having 
the composition 2Fe(NCS) 3 -f 3H 2 0. Solutions of ferric 
thiocyanate become colorless when treated with reducing 
agents which change ferric iron to ferrous iron. 

Exp. 237. — To solutions of a ferrous salt add (a) potassium ferrocy- 
anide, {b) potassium ferricyanide, (c) and ammonium thiocyanate. 

Next add the same reagents to solutions of ferric chloride. Note fully 
the results, and also how ferrous may be distinguished from ferric salts. 

The solution of ferrous salt used should be free from ferric salt, and 
may be prepared by dissolving pure ammonium ferrous sulphate in hot 
water, which has been previously boiled to expel the air, and then cool- 
ing the solution ; or iron wire may be dissolved in hydrochloric acid, 
and the solution poured into boiled water. 



Cobalt, Co. 

Atomic Weight, 59. Density, 8.7. 

Cobalt is not an abundant element, nor are its ores widely 
distributed. The pure metal is prepared by igniting the 
oxide in hydrogen, and the gray powder obtained is melted in 



COBALT. 415 

a lime crucible. Cobalt is somewhat more fusible than iron, 
has the lustre of nickel, is malleable, possesses great tenacity, 
and is magnetic even at a red heat. Impurities render it 
brittle and more fusible. It does not tarnish in common air, 
but oxidizes at a red heat. Electro-deposited cobalt cannot 
be distinguished in appearance from nickel-plate. Little use 
has been made of metallic cobalt because of its high cost, but 
its compounds have long been used for the splendid blue they 
impart to glass. 

Cobalt forms two classes of compounds, viz., cobaltous, 
containing bivalent cobalt, and cobaltic, which exhibit a 
higher valence of cobalt. The simple cobaltic salts are un- 
stable. 

CI 
Cobaltous Chloride, CoCl 2 or Co<« 1 , is obtained in blue 

scales when chlorine acts on warm powdered cobalt. A solu- 
tion is prepared by dissolving the carbonate or any of the ox- 
ides of cobalt in hydrochloric acid. The solution has a light- 
red color, which changes to blue on evaporation ; dilution 
with water restores the red color. 

Cobaltous Oxide, CoO or Co=0, is obtained as an olive-green 
powder by heating the higher oxides in hydrogen to a tem- 
perature not above 300°. 

Cobaltous Hydroxide, Co(OH) 2 . — Potassium hydroxide pro- 
duces in solutions of cobaltous salts a blue precipitate of basic 
salts, which on boiling changes into pale-red cobaltous hydrox- 
ide. It rapidly becomes dark brown, owing to formation of 
cobaltic hydroxide. 

Cobaltous Sulphide, CoS or Co=S, is obtained as a gray crys- 
talline mass by heating cobalt with sulphur. Ammonium sul- 
phide precipitates a black hydrous sulphide which is insoluble 
in dilute acids. The sulphides Co a S 3 , Co 3 S 4 , and CoS a are 
known. 



416 THE EIGHTH GROUP. 

Cobaltous Nitrate, Co(N0 3 ) 2 or ^ 2 Iq>Co, is prepared by 

dissolving the carbonate in nitric acid. The solution when 
slowly evaporated yields red crystals containing six molecules 
of water. The salt is a valuable reagent in blow-pipe analysis. 

Cobaltous Sulphate, CoS0 4 or S0 2 <q>Co. — A solution of 

the oxide or carbonate in sulphuric acid deposits, when evapo- 
rated at ordinary temperature, red crystals of CoS0 4 -f 7H 2 0, 
which have the same form as green vitriol. At 40° to 50° crys- 
tals of CoS0 4 + 6H 2 separate, which are isomorphous with 
the hexhydrated magnesium, zinc, and nickel sulphates. 

Cobaltous Ammonium Sulphate, Co(NH 4 ) 2 (S0 4 ) 2 -f 6H 2 0, 
forms red crystals, which are isomorphous with ferrous am- 
monium sulphate. 

Cobaltous Carbonate, CoC0 3 , and the hydrate, CoC0 3 + 
6H 2 0, are known. Sodium carbonate produces in hot co- 
baltous solutions a bluish or violet precipitate of basic car- 
bonates. 

Cobaltic Oxide, Co 2 3 or 0=Co-0-Co=0, remains as a dark 
powder when the nitrate is gently ignited. 

Cobaltous-Cobaltie Oxide, Co 3 4 or = Co-0-Co-0-Co=0. 

— This oxide is formed when either of the other oxides is ig- 
nited in air. It is analogous in composition to magnetic oxide 
of iron. 

Cobaltic Hydroxide, Co(OH) 3 . — Alkali hypochlorites produce 
in solutions of cobaltous salts a black precipitate of cobaltic 
hydroxide. This hydroxide and cobaltic oxide dissolve in 
cold hydrochloric acid to a brown solution (of CoCl 3 ?), which 
decomposes readily, with evolution of chlorine. 



NICKEL. 417 

Potassium Cobalticyanide, K 3 (C 3 N 3 ) 2 Co, is an analogue of 
potassium ferricyanide, with which it is isomorphous. On 
mixing solutions of potassium cyanide and a cobaltous salt a 
red precipitate of Co(CN) 2 is formed, which dissolves in excess 
of potassium cyanide as Co(CN) 2 .2KCN. When a dilute so- 
lution of this compound containing an excess of potassium 
cyanide and some hydrocyanic acid is boiled potassium co- 
balticyanide is formed : 

Co(CN) 2 .2KCN + KCN + HCN = K 3 (C 3 N 3 ) 2 Co + H. 

Cobalticyanic acid is obtained by decomposing the copper 
salt of it by hydrogen sulphide. 



Nickel^ Ni. 

Atomic Weight, 58. Density, 8.9. 

Nickel is not an abundant element. The pure metal is ob- 
tained by reducing the oxide with hydrogen. It is almost as 
white as silver, is capable of brilliant polish, is malleable and 
ductile. The tenacity of nickel exceeds that of iron. Nickel, 
like iron, combines with carbon and silicon, which render it 
brittle and more fusible. The cast nickel of commerce also 
contains copper, cobalt, and iron. Nickel is attracted by the 
magnet. 

Nickel is not changed by long exposure to air, nor tar- 
nished by hydrogen sulphide, and oxidizes only with difficulty 
at high temperatures. These properties render it a valuable 
and cheap substitute for silver for plating other metals. 
Moreover, electro-nickel plate is harder and resists wear bet- 
ter than silver. Nickel-plating is used to prevent iron and 
steel implements from rusting, and also for ornamenting many 
articles. The bath best adapted for the deposition of nickel 
is a solution saturated at ordinary temperature of nickel am- 



418 THE EIGHTH GBOTJF. 

monium sulphate. To maintain the strength of the solution, 
and also to prevent its becoming acid, plates of cast nickel are 
hung in the bath and connected with the positive electrode, 
the nickel dissolving from the plates as fast as deposited on 
metallic articles in the bath connected with the negative 
electrode. 

Nickel is largely used in alloys, the most important of 
which is German silver. This is an alloy of copper, zinc, and 
nickel, containing 16 to 20 per cent of the latter metal. Alloys 
of nickel are used for smaller coins ; the five-cent pieces of the 
United States coinage consisting of 1 part of nickel and 3 
parts of copper. The same proportion is used in small coins 
in Germany and Belgium. 

Mckel forms but one class of salts, and in these it is bi- 
valent ; the oxide, Ni 2 3 , and the sulphide, NiS 2 , indicate a 
higher valence. The hydrous salts are green, and when de- 
hydrated are usually yellowish. Most solutions of nickel are 
green. 

CI 
Nickel Chloride, NiCl 2 or Ni< c ,. — Finely divided nickel 

burns brilliantly when heated in chlorine, and sublimes in 
pale golden-yellow scales, which dissolve slowly in water con- 
taining hydrochloric acid. Exposed to the air they gradually 
absorb moisture and become green, and are then easily sol- 
uble. A solution of the chloride is best prepared by treating 
the oxide or carbonate with hydrochloric acid. On concen- 
trating the solution crystals of MC1 2 -f- 6H 2 separate, which 
lose their water on heating. The anhydrous chloride absorbs 
ammonia, with the formation of NiCl 2 + 6NH 3 . This com- 
pound is also formed when ammonia is added to a solution of 
the chloride, and separates in blue octahedrons on cooling a 
hot concentrated solution. When nickel is placed in a solu- 
tion of ammonium chloride, and connected Avith the positive 
electrode of a battery, the metal dissolves slowly. Solutions 



NICKEL. 419 

thus prepared, containing nickel chloride, ammonia, and am- 
monium chloride, have been used in nickel-plating. 

Nickel Monoxide, NiO or Ni=0, is obtained as a greenish 
powder when the hydroxide, carbonate, or nitrate of nickel is 
intensely ignited. It does not take up oxygen on prolonged 
heating in air. 

OH 

Nickel Hydroxide, Ni(0H) 2 or Ni< 0H . — Sodium or potas- 
sium hydroxide produces in solutions of nickel salts a light- 
green precipitate of nickel hydroxide. It dissolves in am- 
monia water, forming a solution which has the property of 
dissolving silk. 

Nickel Sesquioxide, N 2 3 or 0=Ni-0-Ni=0, is prepared by 
gently igniting nickel nitrate. It is a black powder, which 
changes into monoxide on heating intensely. It dissolves in 
hydrochloric acid, with evolution of chlorine, and formation 
of NiCl 2 . 

Nickel Trihydroxide, Ni(0H) 3 , is formed when the green 
nickel hydroxide is suspended in a solution of potassium hy- 
droxide through which chlorine is passed : 

Ni(OH) 2 -f KOH + CI = Ni(OH), + KC1. 

It loses water and oxygen on heating, and behaves with hy- 
drochloric acid like the sesquioxide. 

Nickel Nitrate, Ni(N0 3 ) 2 or jxo -0 >Ni ' * s °^ined by dis- 
solving nickel in nitric acid. On evaporating the solution, 
the salt crystallizes with six molecules of water. 

Nickel Cyanide, Ni(CN) 2 . — On adding potassium cyanide to 
a solution of a nickel salt, an apple-green precipitate of nickel 
cyanide is first formed, which dissolves in an excess of pot as- 



•420 THE EIGHTH GROUP. 

sitini cyanide as Xi(CX) 2 .2KCX. This double salt is decom- 
posed by acids, with separation of nickel cyanide. Xickel 
does not form compounds analogous to the ferri- and cobalti- 
cyanides. 

Nickel Sulphate, NiS0 4 or S0 2 <q>Ni.— This salt is pre- 
pared by dissolving the oxide or carbonate in sulphuric acid. 
If the solution is evaporated at a temperature of 15° to 20°, 
green crystals of the hydrated salt XiS0 4 + TH a O are formed, 
which are isomorphous with the corresponding salts of zinc 
and magnesium. At 50° to T0° crystals separate, containing six 
molecules of water. Hydrous nickel sulphate becomes anhy- 
drous when heated above 280°. Xickel sulphate is very sol- 
uble, 100 parts of water at 16° dissolving 37.1 parts, and at 
70° 61.9 parts, of XiS0 4 . The anhydrous sulphate absorbs 
ammonia gas, forming the compound XiS0 4 .6XH 3 , while from 
a solution of the sulphate in concentrated ammonia water 
dark-blue crystals are obtained, having the composition 
XiS0 4 .lXH s + 2H 2 0. 

Ammonium Nickel Sulphate, (NH 4 ) 2 Ni(S0 4 ) 2 + 6H 2 0.— This 
salt is manufactured in large quantities for use in nickel-plat- 
ing by adding ammonium sulphate to a concentrated solution 
of nickel sulphate. The precipitate which is formed is puri- 
fied by recrystallization. At 20° 100 parts of water dissolve 
5.9 parts, and at 85° 28.6 parts of (XHJ 2 Xi(S0 4 ) 2 . 

Nickel Carbonates. — Alkali carbonates throw down from so- 
lutions of nickel salts a pale-green precipitate of basic carbon- 
ates. The metacarbonate XiC0 3 is obtained by heating a 
solution of nickel chloride with calcium carbonate to 150°. 

Nickel Monosulphide, NiS, occurs as the mineral millerite, 
which is sometimes found in hair-like crystals. Xickel mono- 
sulphide is obtained artificially by heating together nickel and 



RUTHENIUM. 421 

sulphur. On adding ammonium sulphide to a solution of a 
nickel salt a black hydrous sulphide is precipitated. 

Nickel Disulphide, NiS 2 , is obtained in the form of a dark- 
gray powder by fusing a mixture of nickel carbonate, sodium 
carbonate, and sulphur, and treating the product with water. 



Ruthenium, Ru. 

Atomic Weight, 103. Density, 12.26. 

Ruthenium is a brittle, white, difficultly fusible metal. Aqua regia 
has little action on it, but molten potassium hydroxide containing nitrate 
attacks it readily. The finely divided metal oxidizes when heated in air. 

Ruthenium Dichloride, RuCl 2 , is formed, together with the trichloride, 
when the pulverized metal is heated in chlorine ; the trichloride vola- 
tilizes, leaving the dichloride as a black powder, insoluble in water and 
acids. 

Ruthenium Trichloride, RuCl 3 . — In order to prepare this chloride an 
acid is added to a solution of potassium ruthenate, and the precipitate 
is dissolved in hydrochloric acid. The solution thus obtained yields 
ruthenium trichloride on evaporation. 

Ruthenium Tetrachloride, RuCl 4 , is obtained by dissolving the tetrahy- 
droxide in hydrochloric acid. It combines with potassium chloride to 
form the salt K 2 RuCl e , an analogue of potassium chloroplatinate. 

Ruthenium Monoxide, RuO, is obtained by heating the dichloride with 
sodium carbonate. It is a dark-gray powder, insoluble in acids. 

Ruthenium Sesquioxide, Ruo0 3 , is formed when the finely divided metal 
is heated in air. It has a dark-blue color, and is insoluble in acids. 

Ruthenium Dioxide, Ru0 2 , is formed when the sulphate is intensely 
heated in air. 

Ruthenium Tetroxide, RuO,, sublimes when chlorine is passed into a 
solution of potassium ruthenate. It melts at 40°, boils at about 100°, and 
at a little higher temperature decomposes with explosive violence. 

Ruthenium Trihydroxide, RuOH) :! , separates as a dark- brown precipitate 
on adding potassium hydroxide to a solution of the trichloride. 

Ruthenium Tetrahydroxide, RuiOHt) -f- 3H.O, is formed when a solution 
of ruthenium bisulphate is treated with potassium hydroxide, and then 
evaporated. 



422 THE EIGHTH GROUP. 

Ruthenium Bisulphate, Ru(S0 4 ) 2 . — Hydrogen sulphide precipitates 
RuS 2 from a solution of ruthenium trichloride. This sulphide is con- 
verted by nitric acid into bisulphate, which remains as a deliquescent 
yellow powder on evaporating the solution. 

Potassium Ruthenate, K 2 Ru0 4 , is formed when finely divided ruthe- 
nium is fused with potassium hydroxide mixed with nitre. Its aqueous 
solution is reddish yellow. 

Potassium Perruthenate, KRu0 4 , is prepared by acting on the pre- 
ceding salt with chlorine : 

K 2 Ru0 4 + Cl = KRu0 4 + KCl. 

From the green solution thus obtained black crystals separate, isomor- 
phous with potassium permanganate. 

Potassium Ruthenocyanide, K 4 (C 3 N 3 ) 2 Ru-J-3H 2 0, is isomorphous with 
potassium ferrocyanide. 



Rhodium, Rh. 

Atomic Weight, 104. Density, 12.1. 

Rhodium is a very hard grayish metal, less ductile and fusible than 
platinum. The metal oxidizes at a red heat, best in form of powder. 
It is not attacked by acids, but if alloyed with lead or platinum it dis- 
solves in aqua regia. 

Rhodium Trichloride, RhCl 3 , is the only known halide of rhodium. It 
is formed when the metal is heated in chlorine. It is insoluble in acids, 
even aqua regia, and only gives up its chlorine at a red heat. 

Rhodium Monoxide, RhO, is obtained as a dark powder when Rh(OH), 
is heated. 

Rhodium Sesquioxide, Rh 2 3 , is obtained by heating the nitrate. 

Rhodium Trihydroxide, Rh(0H) 3 . — When a solution of Ka 3 RhCl 6 is 
heated with a concentrated solution of potassium hydroxide a black 
precipitate of rhodium trihydroxide is formed. It is insoluble in acids. 
If a cold dilute solution of potassium hydroxide is used, another modi- 
fication is obtained, which is soluble in acids and potassium Jiydroxide. 

Rhodium Dioxide, Rh0 2 , is insoluble in acids. It is formed by repeated 
fusion of the finely divided metal with potassium hydroxide and nitrate. 

Rhodium Tetrahydroxide, Rh(0H) 4 , is formed by the prolonged action 
of chlorine on an alkaline solution of the trihydroxide. It dissolves in 



PALLA DIUM. 423 

hydrochloric acid to a blue solution, which slowly gives off chlorine, 
and changes to a red color. 

Sodium Rhodium Chloride, Na 3 RhGl 6 , is prepared by heating a mixture 
of rhodium and common salt in chlorine. It is soluble, and crystallizes 
with 12 molecules of water. The corresponding ammonium salt sepa- 
rates when a solution of rhodium chloride and ammonium chloride 
evaporates spontaneously. 

Rhodium Nitrate is obtained as a gummy mass by evaporating a solu- 
tion of rhodium hydroxide in nitric acid. 

Rhodium Sulphate, Rh 2 (S0 4 ) 3 , separates in hydrous crystals when a so- 
lution of rhodium hydroxide in sulphuric acid is evaporated. 



Palladium,, Pd. 

Atomic Weight, 106. Density, 12. 

Palladium is a white metal resembling platinum and silver 
in lustre and color. It is the most fusible of the platinum 
metals. It welds at a red heat more easily and is softer and 
more malleable than platinum. It volatilizes in the oxyhy- 
drogen flame, giving off green vapors, which condense to a 
dust consisting of oxide and metal. At a red heat the metal 
becomes covered with a blue film of oxide, which disappears at 
higher temperatures. Palladium dissolves in hot nitric acid, 
more readily if it contains nitrous acid. Hydrochloric acid 
dissolves palladium sponge in contact with air. The metal is 
but little used in the arts. On account of its retaining its 
lustre in air it has been used for the graduated surfaces of 
instruments. 

Palladium Hydride. — Palladium possesses greater power of 
absorbing hydrogen than any other metal. At common tem- 
peratures palladium foil absorbs upwards of 370 volumes of 
the gas, and at 100° about 650 volumes. When palladium is 
made the negative electrode in dilute sulphuric acid it takes 
up as much as 960 volumes of hydrogen. The compound 
retains the appearance of the metal, is tough, and more mag- 



424 THE EIGHTH GROUP. 

netic than pure palladium. According to some investigators 
the two elements form the compound Pd 2 H, which takes up 
more hydrogen. Palladium hydride gives up all its hydrogen 
at 100° in vacuum, and also when ignited. 

Palladium forms two classes of compounds, viz., palladious 
and palladic. The simple salts of the common acids are palla- 
dious compounds ; palladic salts doubtless exist in solution, 
but they easily change to palladious salts. 

Palladious Chloride, PdCl 2 . — When a solution of palladium 
in aqua regia is evaporated this compound remains as a brown 
mass. It fuses at a red heat, loses half its chlorine, and 
changes to palladium monochloride, PdCl. 

Palladious Iodide, Pdl 2 , is formed by the direct union of its 
elements. Palladium is blackened when tincture of iodine is 
evaporated upon it. Potassium iodide produces in solutions 
of palladious chloride a black precipitate of palladious iodide ; 
a very characteristic reaction. 

Palladious. Oxide, PdO, remains as a black powder when the 
nitrate is heated. Alkali carbonates precipitate from solu- 
tions of palladious salts a hydroxide, or possibly a basic car- 
bonate, soluble in acids. The precipitate is converted at a 
faint red heat into the oxide Pd 2 0. All the palladium oxides 
are completely decomposed on intense ignition. Sulphides 
analogous to the above oxides are known, also PdS 2 . 

Palladious Nitrate, Pd(N0 3 ),, is prepared by dissolving the 
metal in nitric acid. It is very deliquescent. 

Palladious Sulphate, PdS0 4 , is obtained by dissolving the 
metal in sulphuric acid containing a little nitric acid. It is 
decomposed by water into a basic salt. 

Palladic Oxide, Pd0 2 , is formed when ammonium palladic 
chloride is boiled with potassium hydroxide. It dissolves in 



osmium. 425 

hot dilute hydrochloric acid, with evolution of chlorine, and 
formation of PdCl 2 . 

Palladic Chloride, PdCl 4 , has not been isolated. The aqua 
regia solution of the metal doubtless contains hydrogen pal- 
ladic chloride, H 2 Pd01 6 . 

Potassium Palladic Chloride, K 2 PdCl 6 . — This salt is ob- 
tained when an aqua regia solution of palladium is treated 
with potassium chloride and then evaporated. It is insoluble 
in alcohol and in a solution of potassium chloride. 



Osmium, Os. 

Atomic Weight, 192? Density, 22.5. 

This element is remarkable for its density, which is greater than that 
of any other substance. Metallic osmium is obtained in the form of a 
powder by passing the vapor of the tetroxide mixed with carbon mo- 
noxide and dioxide through a hot porcelain tube. Crystalline osmium 
is formed when amorphous osmium is fused with tin. The latter metal 
is removed by hydrochloric acid, and the crystals are finally heated in 
dry hydrochloric acid gas. The pure metal has a fine blue color, crys- 
tallizes in cubes or rhombohedrons nearly like cubes, which are harder 
than glass. Osmium volatilizes without melting when exposed in a 
lime furnace to the intense heat of the oxy hydrogen flame. Amor- 
phous osmium which has not been intensely ignited is soluble in aqua 
regia and nitric acid. The metal is not used in the arts, but iridosmine, 
a native alloy of iridium and osmium, is valued for points to gold pens. 

Osmium in its highest oxide has a valence of eight, other compounds 
exhibit a lower valence. 

Osmium Trichloride, OsCl 3 , is known only in combination with other 
chlorides, as for example, 3KC1.0sCl 3 + 3H 2 0. 

Osmium Tetrachloride, OsCl,, is obtained in the form of a red subli- 
mate when the metal is heated in dry chlorine. The concentrated 
aqueous solution is yellow, the dilute green. The solution decomposes 
on standing, lower oxides separating, and hydrochloric acid and os- 
mium tetroxide remaining in the solution. 

Potassium Chlorosmate, KjOsCl,,, crystallizes out when a solution con- 



426 THE EIGHTH GKOUP. 

taining potassium chloride, hydrochloric acid, and osmium tetroxide 
is evaporated. It is soluble in water and insoluble in alcohol. 

The Osmium Oxides are OsO, Os 2 3 , Os0 2 , and Os0 4 . 

Osmium Tetrahydroxide, 0s(0H) 4 , separates as a black precipitate when 
alcohol is added to an aqueous solution of the tetroxide. It forms un- 
stable salts. 

Osmic Acid, H 2 0s0 4 , has not been isolated, but a number of its salts 
have been prepared. 

Potassium Osmate, K 2 0s0 4 + 2H 2 0, is obtained by dissolving the tetrox- 
ide in a solution of potassium hydroxide and adding alcohol. It crys- 
tallizes in dark-red octahedrons. 

Osmium Tetroxide, 0s0 4 , is cornmonty known as osmic acid. It is 
formed when the metal is burned in air, or is oxidized by nitric acid. 
The tetroxide is soft, crystalline, and volatile at ordinary temperature. 
It melts at a gentle heat, and boils at 100°. Its gas density has been 
found by experiment to be 128.5; theory requires 131 for Os0 4 . Os- 
mium tetroxide is very poisonous. Its aqueous solution does not redden 
litmus, and is valuable in microscopy for coloring animal tissues. 

Osmiocyanic Acid, H 4 (C 3 N 3 ) 2 0s, separates on addition of hydrochloric 
acid to a solution of its potassium salt. 

Potassium Osmiocyanide, K 4 (C 3 N 3 ) 2 0s -f- 3H 2 0, is an analogue of potas- 
sium ferrocyanide. 






Iridium^ Ir. 

Atomic Weight, 193. Density, 22.4. 

The metal is obtained in a gray porous mass, similar to platinum 
sponge, by heating ammonium iridic chloride. It may be fused in a 
lime furnace with an oxyhydrogen flame, but coal-gas and oxygen will 
not answer. Fused iridium has the lustre of polished steel. It is 
harder than platinum, is only slightly malleable, and breaks with a fine- 
grained fracture. Iridium black separates as an impalpable powder 
when an alcoholic solution of iridium sulphate is exposed to sunlight. 
It is more active in causing the union of combustible gases with oxygen 
than platinum black. Fused iridium is soluble in no acid, but when 
alloyed with much platinum dissolves in aqua regia. Molten alkali hy- 
droxides and potassium disulphate attack iridium, and chlorine com- 
bines with it at a red heat. 

Iridium forms two series of compounds corresponding to iridious 
chloride, IrCl 3 , and iridic chloride, IrCl 4 . The metal appears to be 



IRIDIUM — PLATINUM. 427 

bivalent in IrS and in IrS0 3 , the latter having been obtained in com- 
bination with sodium sulphite and other salts. 

Iridious Chloride, IrCl 3 , sublimes when iridium is heated in chlorine. 
A solution of this chloride is prepared by treating iridic chloride with 
sulphur dioxide. 

Potassium Irido-chloride, K 3 IrCl 6 + 3H 2 0, is obtained by reducing po- 
tassium chloriridate with hydrogen sulphide. The solution after ad- 
dition of potassium chloride yields green crystals on evaporation. 

Iridious Oxide or Iridium Sesquioxide, lr 2 3 , is formed when iridium 
black or sponge is ignited in air. It is completely decomposed at 
temperatures above 1000°. 

Irido- potassium Sulphate, Ir 2 (S0 4 ) 3 + 3K 2 S0 4 . — This double salt remains 
as a green powder when an iridium compound is fused with potassium 
disulphate and the product is treated with a solution of potassium 
disulphate. It is soluble in water and dilute sulphuric acid. 

Iridic Chloride, IrCl 4 . — A solution of iridic hydroxide in hydrochloric 
acid, or of iridium black in aqua regia, probably contains chloriridic 
acid, H 2 IrCl 6 . The solution is said to leave the tetrachloride when 
evaporated at not above 40°. 

Potassium Chloriridate, K 2 IrCl 6 , is obtained by adding potassium chlo- 
ride to a solution of iridic chloride. It is sparingly soluble in cold water 
and insoluble in alcohol. 

Iridic Hydroxide, Ir(0H) 4 . — On addition of an alkali hydroxide to a so- 
lution of iridic chloride a dense blue precipitate of iridic hydroxide is 
formed. 

Iridic Oxide, Ir0 2 , is a heavy black powder obtained by heating iridic 
hydroxide. It is insoluble in acids. 

Iridicyanic Acid, H 3 (C 3 N 3 ) 2 Ir, and Potassium Iridicyanide, K 3 (C 3 N 3 ) 2 Ir, 
have been prepared. 

The Iridium Sulphides are IrS, Ir 2 S 3 , and IrS 2 . 



Platinum, Pt. 

Atomic Weight, 195. Density, 21.5. 

Platinum is almost as white as tin, is moderately hard and 
tenacious, and ranks next to gold and silver in malleability 
and ductility. It may be drawn into fine wire or rolled into 
thin sheets. It is infusible in a wind furnace and in the 
flame of aBunsen burner or blast-lamp; but very tine platinum 



428 THE EIGHTH GROUP. 

wire may be fused in a blow-pipe flame, and also in the under- 
edge of an illuminating-gas jet. 

Platinum is not oxidized when heated in air or oxygen, nor 
is it attacked by any single acid. The most common solvent 
for the metal is aqua regia, in which it is slowly soluble. 
Liquids containing free chlorine also dissolve it. Platinum 
is not affected by molten potassium or sodium carbonates, but 
it is attacked when nitre, the alkali hydroxides, cyanides, or 
sulphides are fused in contact with it. 

Metallic platinum is obtained by igniting ammonium chloro- 
platinate in the form of a gray porous mass known as platinum 
sponge, which may be welded when red-hot into compact 
metal, or fused in a lime crucible by means of a blow-pipe 
flame supplied with coal gas and oxygen. 

When platinum is precipitated from solutions by a reduc- 
ing agent it is obtained as a very fine powder known as 
platinum black. Platinum, like other solids, condenses gases 
on its surface, and it possesses in a marked degree the property 
of causing the union of oxygen with combustible gases. 
Platinum black condenses over 800 times its volume of 
oxygen ; platinum sponge also condeuses the gas, but not in 
such quantity. 

Platinum, on account of its infusibility and permanence 
when in contact with acids, alkaline solutions, or molten 
alkali carbonates, is indispensable to the chemist for crucibles, 
dishes, and other apparatus. Gold also resists the action of 
chemicals, but it costs three to four times as much as platinum, 
and is moreover too fusible for many purposes. Platinum is 
used in the chemical industry for the large stills required in 
the concentration of sulphuric acid, for evaporating dishes, 
and other purposes. An alloy of platinum containing several 
per cent of iridium is harder and less readily attacked by 
various reagents than pure platinum, and such an alloy is 
commonly employed for platinum apparatus. 

Platinum occurs alloyed with small quantities of the plati- 



PLATIKUM. 429 

num metals, and is found in minute traces in some minerals 
and in most silver. Little was known about it until the middle 
of the last century, and the metal did not come into use until the 
present century, when the increased knowledge of chemistry 
made it possible to purify and work it. 

Platinum forms two series of compounds, viz., the platin- 
ous, containing bivalent platinum, as in PtCl 2 ; and the 
platinic compounds, containing tetravalent platinum, as in 
PtCl 4 . 

CI 
Platinous Chloride, PtCl 2 or Pt < «,, is obtained by heating 

chloroplatinic acid to 300°, as a grayish-green powder, which 
leaves metallic platinum on ignition. Platinous chloride is 
insoluble in water, but soluble in hot hydrochloric acid, with 
formation of chloroplatinous acid, H 2 Pt01 4 . A number of 
salts of this acid have been prepared, as, for example, potas- 
sium chloroplatinite, K 2 PtCl 4 , and calcium chloroplatinite, 
CaPtCl 4 + 8H 2 0. 

Platinous Oxide, PtO. — When platinous hydroxide is cau- 
tiously heated the oxide is obtained as a gray powder, which 
is reduced to the metal on ignition. 

Platinous Hydroxide, Pt(0H) 2 , is precipitated when sodium 
hydroxide is added to a solution of potassium chloroplatinite. 
It is a weak base. 




Platinum Tetrachloride, ) p , ri p+ ^ 
Platinic Chloride, j rtUi * or ri 

~Cl 

compound, in the hydrous state, is prepared by adding the 
amount of silver nitrate required to separate one third o\' 
the chlorine from chloroplatinic acid, and evaporating at 
ordinary temperature the solution filtered from the silver 
chloride formed, when large red crystals will be obtained, 
having the composition PtUl 4 + 5II 2 0. This substance loses 



430 THE EIGHTH GKOUP. 

four molecules of water at 100°, but all of the water cannot 
be expelled without decomposing the salt. 

Potassium Platinonitrite, K 2 Pt(N0 2 ) 4 , is an example of a 
series of compounds which do not appear to be double salts. 
It is obtained by acting on potassium chloroplatinite with 
potassium nitrite. Hydrogen sulphide fails to precipitate 
platinum sulphide from its solution. 

Chloroplatinic Acid, H 2 PtCl 6 , crystallizes with six molecules 
of water in brownish-red deliquescent prisms from a solution 
of platinum in aqua regia which has been repeatedly evapo- 
rated with hydrochloric acid to remove all nitric acid. This 
compound is also called platinum or platinic chloride. Chloro- 
platinic acid readily exchanges its hydrogen for metals, form- 
ing a series of salts known as chloroplatinates or platini- 
chlorides. 

Potassium Chloroplatinate, K 2 PtCl 6 , separates on addition 
of potassium hydroxide or a potassium salt to a solution of 
chloroplatinic acid as a yellow crystalline powder. It crystal- 
lizes from hot water in small reddish-yellow octahedrons. 
At ordinary temperature 100 parts of water dissolve about 1 
part, and at 100° 5.18 parts of the salt It is insoluble in 
alcohol, and in a saturated solution of potassium chloride. 

Sodium Chloroplatinate, Na 2 PtCl 6 is easily soluble in water 
and alcohol. It crystallizes with six molecules of water, which 
it loses at 100°. 

Ammonium Chloroplatinate, (NH 4 ) 2 PtCl 6 , is prepared by 
adding ammonium chloride to a solution of chloroplatinic 
acid. It crystallizes in octahedrons similar in appearance to 
the potassium salt. It is less soluble in water than the latter, 
is but slightly soluble in a solution of ammonium chloride, 
and is insoluble in alcohol. 



SUMMARY. 431 

Platinic Hydroxide, Pt(0H) 4 , possesses both basic and acid 
properties, dissolving in dilute acids, and in a solution of 
sodium hydroxide. It is converted into Pt0 2 by gentle heat. 

Platinic Sulphate, Pt(S0 4 ) 2 , is obtained by evaporating a 
solution of the hydroxide or chloride in sulphuric acid. 

Platinous Sulphide, PtS, and Platinic Sulphide, PtS 2 , have 
been prepared. 

Platinum Amines. — There are a large number of very com- 
plex compounds of platinum and ammonia whose hydroxides 
are strong bases which form well-defined salts. The other 
platinum metals also form amine compounds. 

Exp. 238. — a. Dissolve a gram or more of thin scrap platinum in 
warm aqua regia. Evaporate the solution on a water bath to dryness, 
add some hydrochloric acid to the residue, and repeat the evapora- 
tion. Dilute with water, so that each cubic centimeter of solution 
shall contain 0.05 gram of platinum. 

b. To 5 cc of the solution add an excess of a solution of ammonium 
chloride. Collect the ammonium chloroplatinate on a filter, dry, and 
ignite cautiously in a covered porcelain crucible. 

c. Hold a piece of the platinum sponge in a small jet of hydrogen. 
The metal will glow, owing to the combination of oxygen condensed 
from the air and the hydrogen by the metal. The Dobereiner lamp 
consists of a small hydrogen generator for supplying a jet of hydrogen 
which is ignited by platinum sponge. 

d. After using the platinum sponge in the above experiment, heat it 
intensely in the blast-lamp flame. It will not now act so energetically 
when hydrogen comes into contact with it. 



Summary of the Eighth Group. 

This group of elements contains three well-defined sub- 
groups, which fall in three different periods of Mendelejeff's 
classification. The atomic weights, densities, and atomic vol- 



Density. 


Atomic volume 


7.8 


7.2 


8.7 


6.8 


8.9 


6.5 


12.2 


8.4 


12.1 


8.6 


12 


8.8 


22.5 


8.5 


22.4 


8.6 


21.5 


9. 



432 THE EIGHTH GROUP. 

umes of the members of the group are given in the following 
table : 

Atomic weight. 

( Iron, ... 56 

Of period II. \ Cobalt, ... 59 

( Nickel, ... 58 

( Ruthenium, . 103 

Of period III. \ Rhodium, . . 104 

( Palladium, . 106 

( Osmium, . . 192 

Of period V. \ Iridium, . . 193 

( Platinum, . . 195 

The elements of each sub-group differ but little in their 
atomic weights, densities, and atomic volumes. Ruthenium, 
rhodium, and palladium are often called the light platinum 
metals, in distinction from the heavy platinum metals, which 
are osmium, iridium, and platinum. The atomic weights and 
densities of the light platinum metals are little more than half 
the magnitude of the corresponding constants of the heavy 
platinum metals. Accordingly we find that the atomic vol- 
umes of the platinum metals differ but little. The atomic 
weights of iron, cobalt, and nickel bear to those of the light 
platinum metals a relation similar to that stated between the 
two classes of platinum metals. The densities of iron, cobalt, 
and nickel are somewhat higher in comparison with their 
atomic weights, and hence the atomic volumes are lower than 
those of the platinum metals. 

The metals of the group are hard and white, or nearly 
white, osmium excepted, which has a blue color. The mem- 
bers of a sub-group, as might be expected, present analogies in 
chemical properties, and all the nine elements of the eighth 
group are more or less closely related chemically. This is 
partly shown in the following tables of compounds : 



SUMMARY. 

Halides. 



433 



Iron, . 

Cobalt, . 

Nickel, 

Ruthenium, 

Rhodium, 

Palladium, 

Osmium, 

Iridium, 

Platinum, 

Iron, 

Cobalt, . 

Nickel, . 

Ruthenium, 

Rhodium, 

Palladium, 

Osmium, 

Iridium, 

Platinum, 

Iron, 

Cobalt, . 

Nickel, 

Ruthenium, 

Rhodium, 

Palladium, 

Osmium, 

Iridium, 

Platinum 

Iron, . 

Cobalt, . 

Nickel, 

Ruthenium, 

Rhodium, 

Palladium, 

Osmium, 

Iridium, 

Platinum, 



Oxides 
FeO Fe 3 4 
CoO Co 3 4 
NiO 
RuO 
RhO 
PdO 
OsO 



FeCl 2 
CoCl 2 
NiCl 2 
RuCl 2 

PdCl 2 



PtCla 

Fe 2 3 
Co 2 3 
Ni 2 3 
Ru 2 3 
Rh 2 3 



FeCls 
CoCl 3 ? 

RuCl 3 RuCl 4 
RhCl 3 

PdCl 4 
OsCl 3 ? OsCl 4 
. IrCl 3 IrCl 4 

PtCl 4 

Ru0 2 Ru0 4 
Os0 4 



Pd0 2 

Os0 2 

Ir0 2 

Pt0 2 



Os 2 3 
lr 2 3 
PtO 

Oxygen Salts. 
FeS0 4 Fe 2 (S0 4 ) 3 K 2 Fe0 4 

CoS0 4 
NiS0 4 

Ru(S0 4 ) 2 K 2 Ru0 4 KRu0 4 
Rh 2 (S0 4 ) 3 
PdS0 4 

K 2 0s0 4 
Ir 2 (S0 4 ) 3 

PtfS0 4 ) 2 

Complex Cyanides. 

K 4 (0 3 N 3 ) 2 Fe K 3 (C 3 N 3 ) 2 Fe 

Ks(C s N s )aC0 

K 4 (C 3 N 3 ) 2 Ru 



K 4 (C 3 N 3 ) 2 0s 



K 3 iC 3 N 3 ) 2 Ir 



434 THE EIGHTH GEOUP. 

There are other series of complex compounds which have 
not been described in this book, and therefore cannot well be 
used in the study of the group. A more extended discussion 
than space allows would show marked resemblance between 
iron, ruthenium, and osmium ; between cobalt, rhodium, and 
iridium ; and between nickel, palladium, and platinum. 

Iron, cobalt, and nickel possess strongly basic properties, 
forming stable oxygen salts. Ferric salts are more permanent 
than ferrous, while the reverse is the case with cobaltic and 
cobaltous salts ; and nickel forms only nickelous salts. The 
platinum metals are feebly basic, forming unstable oxygen 
salts, and their higher oxides and hydroxides exhibit towards 
strong bases an acidic character. The tetroxides of ruthenium 
and osmium appear to be neutral bodies, possessing neither 
acidic nor basic properties. The compounds of the platinum 
metals with non-metals are more or less readily decomposed by 
heat, while the compounds of iron, cobalt, and nickel resist 
intense ignition, or are reduced to oxides. 






THE ATOMIC THEORY. 



In" the introduction certain physical phenomena have been 
explained by the theory that matter is composed of discrete 
particles called molecules. These are physical units, or 
smallest particles which take part in physical changes. We 
have seen that molecules are divisible into parts, as, for ex- 
ample, those of hydrogen and chlorine, page 61. These 
parts have been termed atoms, and we have taken for granted 
that matter is composed of atoms ; and further, that there is 
one fundamental difference between different kinds of matter 
which cannot be resolved into simpler forms, i.e., a differ- 
ence in the masses of the atoms, or, as commonly stated, a 
difference in atomic weights. 

The quantitative determinations made in our experiments 
have given results in accord with the view that one kind of mat- 
ter always takes part in chemical changes in a definite ratio or 
a simple multiple thereof. Further, we have become familiar 
with a large number of reactions represented by equations in 
which the relative weights of different kinds of matter taking- 
part in the changes are represented. The equations, unless 
hypothetical, are based upon determinations of the quantities 
reacting. 

Philosophers have held different views regarding the divisi- 
bility of matter ; some maintaining that it is continuous and 
capable of infinite division, and others that it is composed of 
particles which cannot be resolved into smaller parts. A 
knowledge of the formation of compounds from unlike sub- 
stances led to the hypothesis that chemical union results from 



436 THE ATOMIC THEORY. 

the combination of unlike particles. At the beginning of the 
present century the particles of different kinds of matter were 
supposed to differ in weight. This was Dalton's hypothesis, 
and is the basis of the atomic theory of the present time. It 
was first suggested to him by his investigations of marsh gas 
and defiant gas, two compounds of carbon and hydrogen, 
which have the following composition, expressed centesimally: 

Marsh gas. Olefiant gas. 

Carbon, 75 85.7 

Hydrogen, .... 25 14.3 

On comparing the results he saw that in marsh gas the 
mass of the hydrogen is one third that of the carbon, and in 
olefiant gas it is one sixth. He then found that other com- 
pounds exhibit similar relations, as for example the oxides of 
carbon and the oxides of nitrogen. The oxides of carbon have 
the composition — 

Carbon monoxide. Carbon dioxide. 

Carbon, 42.86 27.27 

Oxygen, 57.14 72.73 

Here the quantity of oxygen in the first is one and one third 
times that of the carbon, and in the second two and two thirds 
times. 

Such data led Dalton to the 

Law of Multiple Proportions. — Wlien two elements combine 
in different 'proportions, the weight of one element increases 
by simple multiples] the weight of the other element being re- 
garded as constant. 

Dalton sought for an explanation of the fact tnat two ele- 
ments may combine to form different compounds. He sup- 
posed that all substances are made up of indivisible particles, 
called atoms, which cannot be divided, and that compounds 
result from the approximation l1 atoms of different kinds. 



LAW OE DEFINITE PKOPOKTIONS. 43 7 

He therefore concluded that one atom of an element may 
unite with one, two, or more atoms of another element ; and 
since atoms of one kind have a definite weight, the several 
compounds which result from the combination of two ele- 
ments will contain each element in proportion to its atomic 
weight or a simple multiple thereof. 

The atomic theory has developed with the advance of 
science, and much has been learned regarding the various 
properties of atoms. We shall only state those which are of 
fundamental importance in chemistry. According to the 
atomic theory atoms are indivisible ; atoms of different kinds 
differ in mass ; all atoms of one kind of matter possess the 
same mass and identical properties. In other words, the 
atoms of an element are alike in all respects, and the atoms 
of different elements differ in mass, and more or less in 
chemical properties. 

Law of Definite Proportions. — The proportions by weight of 
the constituents of a chemical compound are invariable. 

This law is based on a vast number of analyses and syn- 
theses of well-known compounds, whose composition has been 
found to be constant within the limits of the errors of ex- 
periment. 

Let us consider how the law supports the theory that the 
atoms of an element all have the same weight. For example, 
we will suppose that the atoms of oxygen differ in weight. In 
such case we might expect that sometimes the lighter atoms 
would exceed the heavier in a given quantity of water formed 
synthetically by the union of hydrogen with oxygen. Then 
we should have a sample of water with less than | of its weight 
of oxygen. The numerous syntheses of water have given no 
such result. Since water formed by burning hydrogen with 
varying proportions of oxygen, and by various chemical de- 
compositions, possesses constant composition and properties. 



438 THE ATOMIC THEORY. 

we conclude that the atoms of oxygen composing it do not 
vary from each other in any respect, and that the same is true 
of the atoms of hydrogen. 



Determination of Atomic Weights. 

Before considering the methods for finding atomic weights 
we need to have clearly in mind what an atomic weight repre- 
sents. It is defined on page 47 thus : 

An atomic toeight is a number expressing the ratio of the 
mass of the smallest part of an element entering into combina- 
tion to the mass of an atom of hydrogen, or half the molecule 
of hydrogen. 

We observe that an atomic weight does not represent the 
weight of an atom, but, on the contrary, represents the rela- 
tive mass of an atom compared to the mass of an atom of hy- 
drogen, which is taken as unity. The absolute mass of an 
atom is far too small to be determined by chemical methods; 
but the ratio of the mass of an element entering into combina- 
tion to the mass of hydrogen also entering into combination is 
known, in case of many of the elements, with a very consider- 
able degree of accuracy. The determination of the atomic 
weights of elements which form gasifiable compounds, i.e. 
compounds of which the molecular weights can be determined, 
involves — 

1st. The experimental determination of the ratios of the 
mass of an element entering into the compound molecules to 
the mass of the atom of hydrogen ; and, 

2d. Finding the smallest of these ratios. 

For example, we may take the following compounds of 
chlorine : 



DETERMINATION OF ATOMIC WEIGHTS. 



439 



Hydrogen chloride, 
Methyl chloride. 

Methylene chloride, 

Chloroform, . . . 



Gas 

density. 

. 18.25 
25 25 

. 42.5 



Molecular 
weight. 

36.5 

50.5 

85 



59.75 119.5 



Composition. 
Hydrogen 1, Chlorine 35.5. 
Carbon 12, Hydrogen 3, 

Chlorine 35.5. 
Carbon 12, Hydrogen 2, 

Chlorine 71. 
Carbon 12, Hydrogen 1, 
Chlorine 106.5. 



Here the first column gives the observed gas density, the 
second the molecular weight as deduced by the law of Avoga- 
dro, and the third the composition derived from analysis. 
Evidently the smallest ratio of the mass of chlorine in the mole- 
cules of the above compounds to the mass of the atom of hy- 
drogen is 35.5 : 1. Many other compounds of chlorine are 
known, and in none is a smaller ratio found. Hence we con- 
clude that 35.5 is the maximum atomic weight of chlorine. 
If compounds of chlorine are discovered in which the ratio 
is less than now known, then 35.5 will be proved to be a mul- 
tiple of the true atomic weight. 

For the atomic weight of oxygen we have — 



Gas 
density. 



Water, 9 

Methyl alcohol, . . 16 

Ethyl alcohol, ... 23 

Acetic acid, .... 30 



Molecular 
weight. 

18 



32 



46 



60 



Composition. 
Hydrogen 2, Oxygen 16. 
Hydrogen 4, Carbon 12, 

Oxygen 16. 
Hydrogen 6, Carbon 24, 

Oxygen 16. 
Hydrogen 4, Carbon 24, 

Oxygen 32. 

The smallest ratio of the mass of oxygen in these compounds 
to the mass of the atom of hydrogen is 16:1. From a con- 
sideration of the carbon compounds in the foregoing tables 
we conclude that 12 is the atomic weight of carbon. 

To illustrate further, we will take the compounds of carbon 
and oxygen : 

Molecular 
weight. Composition. 

28 Carbon 12. Oxygon 16. 

44 Carbon 12, Oxygen 32. 





Gas 
density, 


Carbon monoxide, . 


. 14 


Carbon dioxide, 


22 



440 THE ATOMIC THEORY. 

Accordingly, we have carbon 12 and oxygen 16, numbers 
identical with those deduced from the hydrogen compounds 
of these elements. The atomic weights of chlorine, oxygen, 
and carbon are founded upon the composition and molecular 
weights of a very large number of their compounds, but the 
foregoing examples suffice for illustration. 

A number of the elements do not form gasifiable compounds, 
hence the molecular weights of their compounds are unknown. 
The relative masses, which exist in the compounds of such an 
element with elements whose atomic weights are known, are 
found by analysis and synthesis, and the number which is the 
true atomic weight is found by the 

Law of Dulong and Petit. — The specific heat of an element 
in the solid state multiplied by its atomic weight is nearly 
constant. 

For example : 

Specific heat. Atomic weight. Sp. ht. X at. wt. 

Aluminum, 0.225 27 6.1 

Iron, 0.114 56 6.4 

Lead, 0.0314 207 6.5 

Magnesium 0.25 24 6.0 

To illustrate the application of this law in fixing an atomic 
weight, let us take as an example calcium. It was found 
that 20 weights of this metal combine with 35.5 weights of 
chlorine, with 8 weights of oxygen, and displace 1 weight of 
hydrogen from acids in forming salts. All the reactions of 
calcium can be accounted for by assuming that 20 is its com- 
bining weight. It is easy to see that half or double this num- 
ber will answer equally well. Bnnsen found the specific heat 
of metallic calcium to be 0.17. Therefore, since of these pos- 
sible values for the atomic weight 40 is the only one which, 
multiplied by the specific heat, will give nearly 6, we conclude 
that the atomic weight of calcium is 40. 



LAW OF DULONG AND PETIT. 441 

Let us next consider the atomic weights of elements, for- 
tunately few, which do not form gasifiable compounds, and 
whose specific heats are doubtful or unknown. In such case 
we compare the compounds of the elements, in question with 
those of closely related elements whose atomic weights are 
known. For example, calcium, strontium, and barium are 
very similar in chemical deportment, forming compounds 
which possess remarkable similarity in properties. They form 
a clearly-defined group of elements, and in all probability 
their compounds have an analogous composition. Strontium 
chloride contains 43.75 weights of strontium and 35.5 weights 
of chlorine. Its composition may therefore be represented by 
the formula SrCl. But the formula of calcium chloride is 
Ca01 2 , and from analogy for the corresponding strontium salt 
we have SrCl 2 , which requires 87.5 weights of strontium to 
twice 35.5 weights of chlorine. In the same way we may 
compare other classes of calcium and strontium compounds, 
and we conclude that 87.5 is the probable atomic weight of 
strontium. ^The atomic weight of barium rests upon the 
same basis as that of strontium. 

We have seen that atomic weights of some elements are not 
directly compared in the experimental data to the mass of the 
atom of hydrogen, but to the mass of some other element 
whose relation to the atom of hydrogen is known. In fact, 
most of the atomic weights are based upon such indirect com- 
parisons. 

In concluding the discussion, let us endeavor to obtain a 
clear conception of the data required to fix an atomic weight. 

1. The proportion by weight of the element must be de- 
termined with the greatest possible accuracy by analysis or 
synthesis, or both, in compounds which can best be prepared 
in greatest purity. Such compounds should also be of the 
element in question with elements whose atomic weights are 
most accurately known. 



442 THE ATOMIC THEORY. 

2. The determination of the gas density, if possible, of one 
or more compounds of the element. 

3. The determination of the specific heat of the element in 
the solid state. 



Do Atoms Exist ? — This may be regarded as a fundamental 
question in chemistry and physics. The present science of 
chemistry is based upon the atomic theory. Without it we 
have no explanation of the law of definite proportions, the 
law of multiple proportions, and the fact that isomeric com- 
pounds exist. We believe that chemical changes are due to 
combinations of atoms, the separation of groups of atoms, and 
also to the rearrangement of atoms within a molecule. If, 
then, the atomic theory is in accord with all facts and laws of 
physical and chemical science, we must conclude that atoms 
exist. It is often stated that atoms are so small that it 
is not possible to prove their existence. Until, however, facts 
are discovered which are clearly at variance with the atomic 
theory it must be regarded as established. 



THE PEKIODIC LAW. 



The study of successive groups of elements has made evi- 
dent the arrangement in Men dele jeff ^s classification. The 
table, p. 49, presents the elements in order of increasing 
atomic weights, and in groups containing elements more 
or less closely related in properties. The members of a 
group are also arranged in order of their increasing atomic 
weights ; e.g., the first group containing Li, 7 ; Na, 23 ; 
K, 39; Cu, 63.3; Eb, 85.5; Ag, 107.9; Cs, 133; and 
Au, 196. The remarkable fact presents itself that the atomic 
weight of any member of this group is approximately a numeri- 
cal mean of the next lower and next higher atomic weight, 
omitting the last ; thus : 

7 + 39 = ^ 23 + 63.3 = ^ 39 + 85,5 = ^ 



63.3 + 107.9 = 85 ^ 85^+133 = m3> 



Such relations between the atomic weights of the elements of 
a natural group have long been known. 

It has been shown in the summaries of the groups that the 
properties of the members of a group vary with more or loss 
regularity with the atomic weights. 



444 THE PERIODIC LAW. 

Let us now consider some of the simplest relations which 
appear when the elements are arranged according to increas- 
ing atomic weights, and take for this purpose the following 
table : 



Li= 7 


LiCl 


Li 2 


Na=23 


NaCl 


Na 2 


Be= 9 


BeCl 2 


BeO 


Mg=24 


MgCl 2 


MgO 


B =11 


BC1 3 


BA 


Al =27 


A1C1 3 


Ai 2 o a 


=12 


CH 4 


co 2 


Si =28 


SiH 4 


Si0 3 


N =14 


NH 3 


N 2 5 


P =31 


PH 3 


PA 


=16 


OH 2 




S =32 


SH 2 


so 3 


F =19 


FH 




CI =35.5 


C1H 





The capacity to combine with chlorine or hydrogen increases 
regularly from lithium to carbon, and then diminishes regu- 
larly to fluorine. Following fluorine is sodium, which, like 
lithium, combines with one atom of chlorine, and also forms 
an oxide analogous to lithium oxide. From sodium to silicon 
the number of fixed atoms of hydrogen or chlorine increases 
regularly, and then diminishes from silicon to chlorine. It is 
obvious that in this respect there is complete analogy between 
the elements from lithium to fluorine compared with those 
from sodium to chlorine. 

The oxygen compounds exhibit a regularly increasing ca- 
pacity of the elements for oxygen from lithium to nitrogen, 
and again from sodium to sulphur. The first four oxides in 
each column are analogous to the first four chlorides or hy- 
drides, one atom of oxygen being equivalent to two of chlorine 
or of hydrogen. Such of the remaining elements in the table 
as form oxides possess a greater capacity for oxygen than for 
hydrogen or chlorine. 

When we compare the hydroxides, a gradation in properties 
is also striking. The hydroxides which form the most familiar 
salts answer our purpose best. They are 



THE PERIODIC LAW. 445 

LiOH NaOH 

Be(OH) a Mg(OH) 2 

B(OH), Al(OH), 

C0(0H) 2 SiO(OH) 2 and Si(0H) 4 

N0 2 OH P0 2 OH and PO(OH) 3 

S0 2 (OH) 2 

— C10 3 OH 

Lithium hydroxide is a strong base ; beryllium hydroxide 
is also a base, but may be separated from salts by lithium hy- 
droxide ; boron hydroxide possesses no basic properties, but is 
a feeble acid, displaced by even carbonic acid ; carbon hydrox- 
ide is never basic, and is set free by the common strong acids ; 
the hydroxide of nitrogen is nitric acid ; fluorine forms no 
hydroxide or oxy-acid. Following next is sodium hydrox- 
ide, a strong base like lithium hydroxide ; magnesium hy- 
droxide is a weaker base ; aluminum hydroxide is a still 
weaker base, and also possessing slightly acidic properties ; 
silicon hydroxide is never basic, and is displaced by stronger 
acids ; and the hydroxides of phosphorus, sulphur, and chlo- 
rine are powerful acids. 

The elements from lithium to fluorine exhibit a regular 
gradation from strongly basic to strongly acidic characteris- 
tics ; and the same holds good from sodium to chlorine. The 
elements we have considered belong to the first period of 
Mendelejeff 's classification. A careful study of the members 
of the second and third periods will make evident a similar 
variation in properties from element to element, with a return 
to very like properties at certain fixed points. This leads 
to the hypothesis that the chemical properties of an element 
are very largely determined by its atomic weight. It may be 
stated that physical properties, e.g., melting and boiling- 
points, densities, and atomic volumes, bear a relation to, and 
to a greater or less degree are dependent upon, the atomic 
weights. 



446 THE PERIODIC LAW, 

The Periodic Law. — The properties of the elements vary 
from member to member in order of increasing atomic weights, 
and return to more or less closely related properties at fixed 
points in the whole series of the elements; and certain proper- 
appear periodically. 



The classification makes evident the meaning of this law. 
Periodicity appears in the elements of any group, e.g., in case 
of the closely allied metals of the first group, or the halogens 
of the seventh group. Further, we see that in regard to 
valence there is an increase from the first to the seventh 
group, and in a less marked degree to the eighth group. Since 
the valence which an element exhibits depends upon the com- 
pound containing it, and is a variable property, it must be 
borne in mind that while there is evidently a periodicity in 
valence, there are some exceptions to the rule that the high- 
est valence of an element corresponds to the number of the 
group in which the element falls. This is, to a degree, con- 
nected with the fact that members of one group possess certain 
properties in common with some members of another group. 
Lithium in the insolubility of its carbonate and phosphate 
resembles magnesium. Copper in its more stable compounds 
is bivalent, and its more common salts are analogous to the 
salts of the members of the second group. Gold in its tri- 
valent character is allied to the third group. Thallium in 
thallous compounds resembles the alkali metals. Lead, which 
is tetravalent, forms salts analogous in formula to the salts of 
the members of the second group ; its sulphate is insoluble 
like the sulphates of calcium, strontium, and barium. The 
members of the fifth group in their trivaient character re- 
semble the elements of the third group ; iron in ferrous 
compounds is related to the second group, and in ferric com- 
pounds to the third group. Numerous other examples might 
be stated. 



THE PERIODIC LAW. 447 

The element having the lowest atomic weight stands apart 
from the other members of a group. This has been shown in 
the summaries of the groups, and need not here be illustrated by 
examples. In a group whose members falling in period I. are 
non-metals we find with increasing atomic weights that metal- 
lic and basic properties increase. For instance, phosphorus is 
a non-metal, and is only acidic ; arsenic is somewhat metallic in 
character, and feebly basic ; and in antimony metallic and 
basic characters are more pronounced. 



INDEX. 



Absorption of gases by charcoal, 

318 
Acetates, 362 
Acetic acid, 359 
Acid anhydrides, 165 
Acid salts, 165 
Acids, 162 
iEriform bodies, 6 
Agate, 366 
Air, 249 

density of, 10 

weight of, 251 
Alcohol, 357 
Alcohol radicals, 356 
Aldehyde, 358 
Aldehydes, 351 
Alkali-earth metals, 196, 335 
Alkalies, caustic, 137 
Alkali metals, 83 

oxides and hydroxides of, 137 
Allotropy, 111 
Alumina, 299 
Aluminates, 300 
Aluminum, 296 

alloy with copper, 297 

alloy with iron, 297, 298 

bronze, 297 

chloride, 298 

hydroxides, 299 

oxide, 299 

sodium fluoride, 299 

sulphate, 301 

valence of, 298 
Alums, 302 
Amalgam of gold, 104 
Amalgams, 216 
Ametbyst, 366 
Amides, 353 



Amines, 353 
Ammonia, 223 

compounds with metallic salts, 
231 

water, 224 
Ammonias, compound, 352 
Ammonium, 224, 231 

alum, 302 

bromide, 230 

carbonates, 332, 333 

chloride, 228 

chloroplatinate, 430 

fluoride, 229 

hydroxide, 224 

iodide, 230 

magnesium phosphate, 267 

molybdate, 186 

nickel sulphate, 420 

nitrate, 237 

phosphomolybdate, 187 

salts, constitution of, 231 

sulphate, 230 

sulphides, 230 

sodium phosphate, 266 

stannic chloride, S82 

sulpharsenite, 275 

thiocyanate, 345 
Amorphous substances, 12 
Analysis, 64 
Anatase, 373 
Anhydrides, 165 
Anhydrite, 202 
Animal charcoal, 320 
Antifriction metal, 277 
Antimonates, 280 
Antimonie acids, 279 

anhydride, 219 
Antimonous acid, 279 

449 



450 IKDEX. 



Antimonous oxide, 278 
Antimony, 276 

alloys of, 276 

hydride, 277 

chlorides, 277, 278 

oxides and hydroxides of, 278 

pentasulphide, 280 

potassium tartrate, 279 

trisulphide, 280 
Antimonyl chloride, 278 
Aqua ammonia, 224 
Aqua regia, 248 
Archimedes, principle of, 10 
Arragonite, 333 
Argentic chloride, 100 

oxide, 142 
Argentite, 153 
Argentous oxide, 142 
Arsenic, 268 

acid, 275 

anhydride, 274 

di-iodide, 272 

disulphide, 275 

oxides and hydroxides of, 272 

pentasulphide, 275 

pentoxide, 274 

trichloride, 272 

trifluoride, 272 

tri-iodide, 272 

trioxide, 273 

trisulphide, 275 
Arsenious acid, 273 

anhydride, 273 

oxide, 273 
Arseniuretted hydrogen, 270 
Arsine, 270 

Atmo, definition of, 32 
Atmosphere, 249 
Atomic theory, 435 

volume, 84 

weight, definition of, 47 

weights, data for, 441 

weights, determination of, 438 

weights, relations of, 443 

weights, table of, 48 
Atoms, 46, 435, 442 
Auric chloride, 105 

hydroxide, 142 

nitrate, 238 

oxide, 143 
Auroso-auric chloride, 105 
Aurous chloride, 105 






Aurous oxide, 142 
Auryl hydroxide, 143 

nitrate, 238 
Avogadro, law of, 38 
B axing powders, 330 

soda, 329 
Balance, 7 
Barite, 203 
Barium, 203 

chloride, 203 

carbonate, 334 

dioxide, 204 

ferrate, 408 

hydrosulphide, 205 

hydroxide, 204 

hypophosphite, 261 

iodate, 205 

nitrate, 239 

oxide, 204 

sulphate, 205 

sulphides, 205 
Barometer, 31 
Bases, 162 
Basic salts, 165 
Beryllium and compounds of, 195 
Bicarbonate of potash, 332 
Bicarbonate of soda, 329 
Bismuth, 281 

chlorides, 282 

fusible alloys, 281 

oxides, 282, 283 

nitrate, 282 

nitrate, basic, 282 

sulphate, 283 

sulphide, 283 
Bismuthic acid, 283 
Bismuthyl sulphate, 283 
Black-lead, 315 
Bleaching powder, 200 
Blue vitriol, 174 
Blue-stone, 174 
Boiling points, table of, 29 
Bone-black, 320 
Borax, 295 
Boric acids, 293, 294 

oxide, 293 
Boron, 291 

chloride, 292 

fluoride, 292 

hydride, 292 

oxychloride, 293 
Brass, 210 



INDEX. 



451 



Brick, 371 
Britannia metal, 276 
Bloodstone, 366 
Boyle, law of, 6, 33 
Bromic acid, 127 
Bromine, 70 
Brookite, 373 
Brucite, 208 

Bunsen's ice calorimeter, 30 
Burnettizing process, 211 
Butter of antimony, 277 
Cesium, 89 

chloride, 89 

hydroxide, 140 

sulphate, 174 
Cadmium, 213 

chloride, 213 

hydroxide, 214 

iodide, 213 

nitrate, 239 

oxide, 213 

sulphate, 214 

sulphide, 214 
Calamine, 209 
Calcite, 333 
Calcium, 197 

acetate, 362 

carbonate, 333 

chloride, 197 

bromide, 198 

dioxide, 200 

fluoride, 198 

hydroxide, 199 

hydroxychloride, 198 

hypochlorite, 200 

iodide, 198 

nitrate, 238 

oxide, 198 

phosphates, 266 

sulphate, 202 
Calomel, 217 
Calorie, 27 
Calorimeter, 30 
Carbamic acid, 340 
Carbamide, 340 
Carbon, 313 

dioxide, 323 

disulphide, 339 

monoxide, 336 

tests for, 326 
Carbonated water, 325 
Carbonates, 327 



Carbonic acid, 327 

anhydride, 323 

oxide, 336 
Carbonyl chloride, 339 

sulphide, 340 
Carnelian, 366 
Cassiterite, 382 
Cast iron, manufacture of, 396 

varieties of, 396 
Caustic potash, 139 
Caustic soda, 138 
Celestite, 202 

Cerium and compounds of, 375 
Cerusite, 391 
Chalcedony, 366 
Chalcocite, 95 
Charcoal, 317 

absorption of gases by, 318 

animal, 320 

composition of, 318 
Chlorauric acid, 106 
Chloric acid, 125, 133 
Chloride of lime, 200 
Chlorine, 56, 64 

monoxide, 123 

peroxide, 124, 133 
Chloroform, 348 
Chloroplatinic acid, 430 
Chlorostannic acid, 382 
Chlorous acid, 124 
Chrome alum, 179 
Chromic acid, 181 

chloride, 179 

hydroxides, 179 

oxide, 178 

sulphate, 179 
Chromite, 177 
Chromium, 177 

hexfiuoride, 180 

oxychloride, 184 

sesquioxide, 178 

trioxide, 180 
Chromous chloride, 177 

hydroxide, 178 

sulphate, 178 
Chromyl chloride, 184 
Cinnabar, 214 

Classification of elements, 49 
(lav, 370 
Coal, 320 

composition of, 321 
Cobalt, 414 



452 



INDEX. 



Cobalt, derivation of the term, 269 
Cobaltic oxide, 416 

hydroxide, 416 
Cobaltous carbonate, 416 

chloride, 415 

hydroxide, 415 

nitrate, 416 

oxide, 415 

sulphate, 416 

sulphide, 415 
Coke, 316 
Combustion, 110 
Common salt, 86 
Compound ammonias, 352 
Compound radicals, 162 
Copper, 94 

acetate, 362 

arsenite, 274 

cyanides, 343 

dioxide, 142 

nitrate, 237 

pyrites, 94 

sulphate, 174 

tetrantoxide, 140 
Corrosive sublimate, 219 
Cream of tartar, 331 
Cryolite, 299 
Crystallography, 12 

systems of, 14 
Cupric bromide, 97 

chloride, 97 

hydroxide, 141 

iodide, 97 

nitrate, 237 

oxide, 141 

sulphate, 174 

sulphide, 153 
Cuprous bromide, 96 

chloride, 96 

iodide, 96 

oxide, 140 

sulphide, 153 
Cyanic acid, 344 
Cyanogen chloride, 344 
Cyanogen compounds, 341 

gas, 341 
Cyanuric acid, 345 

chloride, 344 
Dalton's hypothesis, 436 
Definite proportions, law of, 437 
Density, 2 
Densities, determination of, 9, 41 



286 



Derived units, 2 

Diamide, 226 

Diamond, 314 

Dicarbon compounds, 354 

Dichromates, 182 

Didymium, and compounds of 

Difference of temperature, 24 

Diffusion, 39 

Dimercurous ammonium chloride, 

232 
Dimethylamine, 354 
Dimorphism, 22 
Diphosphoric acid, 268 
Distilled water, 115 
Disulphuric acid, 169 
Dithionic acid, 170 
Dolomite, 206 

Dulong and Petit, law of, 440 
Dyne, 3 

Earthen ware, 371 
Eighth group, 394 

summary of, 431 
Elements, 46 

classification of, 49 
Emery, 299 
Epsom salt, 208 

Erbium, and compounds of, 310 
Ethane, 355 
Ethyl alcohol, 357 

bromide, 356 

chloride, 355 

iodide, 356 
Ferric acid, 408 

ammonium sulphate, 408 

chloride, 405 

ferrocyanide, 413 

hydroxides, 406 

oxide, 405 

nitrate, 407 

sulphate, 407 

thiocyanate, 414 
Ferricyanic acid, 412 
Ferrocyanic acid, 412 
Ferrous ammonium sulphate, 402 

carbonate, 403 

chloride, 400 

ferric oxide, 406 

ferricyanide, 413 

iodide, 401 

hydroxide, 401 

oxide, 401 

nitrate, 401 



IKDEX. 



453 



Ferrous sulphate, 402 

sulphide, 403 

thiocyanate, 414 
Fifth group, 222 

summary of, 288 
First group, 83 

oxides and hydroxides of, 137 

sulphates of, 172 

sulphides and hydrosulphides 
of, 152 
Flint, 366 
Fluids, 5 
Fluorine, 75 
Force, 3 

Formaldehyde, 348, 351 
Formic acid, 349, 351 
Forms of matter, 5 
Formulas, 47 
Fourth group, 313 

summary of, 392 
Fowler's solution, 274 
Fundamental conceptions, 1 

units, 2 
Fusion, 28 

latent heat of, 29 
Galena, 390 

Gallium and compounds of, 303 
Gas carbon, 316 

Gas densities, determination of, 41 
Gases, 6 

kinetic theory of, 34 

laws of, 33 

measurement of, 32 

pressure and volume of, 31 
Gaseous pressure, explanation of, 

35 
Gay-Lussac, law of, 33 
Germanium, and compounds of, 

377 
German silver, 418 
Gilding, 104 

Glacial phosphoric acid, 268 
Glass, 369 
Glauber's salt, 172 
Gold, 101 

alloys, 102 

amalgam, 104 

coins, 102 

dichloride, 105 

purple, 104 

refining of, 103 

soluble in selenic acid, 171 



Gold sulphate, 176 

sulphides, 154 

trichloride, 105 
Gram weight, 4 
Graphite, 315 
Green vitriol, 402 
Gunpowder, 236 
Gypsum, 202 

Halogens, summaries of, 77, 135 
Heat, 24, 27 

specific, 28 

unit of, 27 
Heavy spar, 205 
Hematite, 394, 406 
Hexagonal system, 20 
Hydraulic cement, 371 
Hydrazine, 226 
Hydriodic acid, 72 
Hydrobromic acid, 71 
Hydrochloric acid, 66 

analysis of, 59 

synthesis of, 58, 60 
Hydrocyanic acid, 342 
Hydrofluoric acid, 76 
Hydrofluosilicic acid, 365 
Hydrogen, 50 

ammonium carbonate, 333 

auric chloride, 106 

auryl sulphate, 176 

bromide, 71 

chloride, 66 

diffusion of, 39 

dioxide, 119 

dioxide, synthesis of, 120, 122 

disulphide, 151 

fluoride, 76 

iodide, 72 

magnesium carbonate, 335 

potassium carbonate, 332 

potassium disulphate, 174 

potassium fluoride, 89 

potassium sulphate ,174 

potassium tartrate, 331 

selenide, 151 

silicon fluoride, 365 

sodium carbonate. 309 

sodium sulphate, 173 

sodium sulphite. 158 

stannic chloride, 382 

sulphide, 148 

telluride, 151 
Hydroxides, definition of, 120 



454 



INDEX. 



Hydroxyl, 120 
Hydroxylamine, 226 
Hypobronious acid, 127 
Hypochlorous acid, 123 

oxide, 123 
Hyponitrous acid, 241 
Hypophosphoric acid, 262 
Hypophospkorous acid, 261 
Hyposulphite of soda, 170 
Hyposulphurous acid, 170 
Indium and compounds of, 304 
Infusorial earth, 367 
Iodic acid, 128 

anhydride, 128 
Iodine, 71 

monoehloride, 74 

pentoxide, 128 

trichloride, 74 

trioxide, 127 
Iridium and compounds of, 426 
Iron, 394 

alum, 408 

cyanides, 408 

disulphide, 404 

magnetic oxide of, 406 

preparation of pure, 398 

properties of, 399 

pyrites, 404 

sesquioxide of, 405 

table of kinds, 395 
Isomerism, 322 
Isometric system, 15 
Isomorphism, 22 
Jasper, 366 
Kaolin, 370 
Kinetic theory, 34 
Kieserite, 208 
Lamp-black, 316 
Lanthanum and compounds of, 

309 
Latent heat of fusion, 29 
Laughing-gas, 242 
Law of Avogad.ro, 38 

Boyle, 6, 33 

definite proportions, 437 

Dulong and Petit, 440 

Gay-Lussac, 33 

multiple proportions, 436 
Laws of gases, 33 
Lead, 385 

alloys of, 387 

acetate, 390 



Lead carbonate, 391 

chloride, 387 

chromate, 181 

dioxide, 389 

ferricyanide, 412 

hydroxy chloride, 387 

in drinking water, 386 

iodide, 388 

oxide, 388 

nitrate, 390 

oxy hydroxide, 389 

suboxide, 388 

sulphate, 390 

sulphide, 390 

tetrachloride, 388 
Lime, 198 
Limestone, 333 
Limonite, 394 

Liquid lrydrogen phosphide, 
Liquids, 5 
Litharge, 388 
Lithium, 84 

chloride, 85 

hydroxide, 137 

oxide, 137 

sulphate, 174 
Lodestones, 406 
Lunar caustic, 238 
Magnesia, 208 
Magnesia alba, 335 
Magnesite, 335 
Magnesium, 206 

carbonate, 335 

chloride, 207 

hydroxide, 208 

nitrate, 239 

oxide, 208 

pyrophosphate, 268 

sulphate, 208 
Magnetite, 394, 406 
Manganese, 78 

black oxide of, 129 

dichloride, 78 

dioxide, 129 

heptoxide, 131 

monoxide, 128 

tetrachloride. 79 

tetrafluoride, 79 

trioxide, 131 
Manganic acid, 131 

oxide, 129 
Manganite, 129 



257 



INDEX. 



455 



Manganous chloride, 78 

hydroxide, 128 

manganic oxide, 129 

oxide, 128 
Marble, 333 
Marsh gas, 346 
Massicot, 388 
Matter, 1 

forms of, 5 

molecular structure of, 22 
Mendelejeff 'stable of the elements, 

49 
Measurement of gases, 32 
Melting points, table of, 29 
Mercuric ammonium chloride, 230 

chloride, 219 

cyanide, 341 

nitrate, 240 

oxide, 219 

orthosulphate, 220 

sulphate, 220 

sulphide, 220 
Mercurous ammonium chloride, 

chloride, 217 

fluoride, 218 

nitrate, 240 

oxide, 218^ 

sulphate, 218 
Mercury, 214 

Metaphosphoric acid, 263, 269 
Metastannic acid, 383 
Methane, 346, 350 

chlorine substitution products 
of, 347 

constitution of derivatives of, 
349 
Methyl alcohol, 348 

amine, 353 

chloride. 347 

iodide, 348 
Microsmic salt, 266 
Millerite, 420 
Mineral coal, 320 
Minium, 389 
Mispickel, 269 
Molecular structure of matter, 22 

velocities, 37 

weights, 38, 62, 63 
Molecules, 22, 35, 40, 61 

size and number of, 40 

velocities of, 38 



Molybdenum and compounds of, 

185 
Monoclinic system, 21 
Mosaic gold, 384 
Motion, 3 

Multiple proportions, law of, 436 
Nickel, 417 

alloys of , 418 

carbonates, 420 

chloride, 418 

cyanide, 419 

hydroxides, 419 

nitrate, 419 

oxides, 419 

plating, 417 

sulphate, 420 

sulphides, 420, 421 
Niobium and compounds of, 285 
Nitrates, 236 
Nitre, 236 
Nitrides, 223 
Nitrites, 241 
Nitric acid, 232 

action of metals on, 234 

action on organic substances, 
233 

anhydride, 247 

oxide, 244 

red fuming, 233 

reduction to ammonia, 234 
Nitrogen, 222 

chloride, 227 

constitution of oxygen com- 
pounds of, 248 

dioxide, 246 

fluoride, 227 

iodide, 227 

oxides and hydroxides of, 232 

pentoxide, 247 

tetroxide, 246 

trioxide, 246 
Nitrous acid, 240 

anhydride, 246 

oxide, 241 
Nitro-hydrochloric acid, 248 
Nitrosyl chloride, 247 
Nitroxyl chloride, 248 
Normal salts, 165 
Oil of vitriol, 159 
Onyx, 366 
Opal, 366 
Orpiment, 269 



456 



LtfDEX. 



Organic chemistry, definition of, 
322 

compounds, 322 
Osmium and compounds of, 425 
Orthorhombic system, 18 
Oxidation, 110 
Oxides, 110 
Oxygen, 107 
Ozone, 110 

Palladium and compounds of, 423 
Paracyanogen, 342 
Paraldehyde, 358 
Paris green, 274 
Pentathionic acid, 170 
Perchloric acid, 126, 133 
Perchromic acid, 183 
Periodic acid, 128, 133 
Periodic law, 446 

system, 49 
Permanganic acid, 131, 134 

oxy chloride, 131 
Pernitric acid, 247 
Petrified wood, 366 
Pewter, 276 
Phosphates, 265 
Phosphine, 256 
Phosphomolybdic acid, 186 
Pbosphonium iodide, 257 
Phosphoretted hydrogen, 256 
Phosphoric acid, di-, 268 

meta-, 268 

ortho-, 264 
Phosphoric acids, constitution of, 
262 

reactions of, 264 
Phosphoric anhydride, 260 
Phosphorous acid, 261 
Phosphorus, 252 

amorphous, 254 

di-iodide, 259 

octahedral, 253 

oxides and hydroxides of, 260 

oxychloride, 259 

pentachloride, 258 

pentafluoride, 258 

pentoxide, 260 

red, 254 

rhombohedral, 255 

trichloride, 258 

trifluoride, 258 

tri-iodide, 259 
Phosphoryl chloride, 259 



Physical magnitudes, 1 
Plaster of Paris, 202 
Platinic chloride, 429 

hydroxide, 431 

sulphate, 431 

sulphide, 431 
Platinous chloride, 429 

hydroxide, 429 

oxide, 429 

sulphide, 431 
Platinum, 427 

amines, 431 

tetrachloride, 429 
Pleomorphism, 22 
Plumbago, 315 
Polychromates, 183 
Polymerism, 322 
Porcelain, 371 
Potash, 331 

nitre, 236 
Potassium, 87 

acetate, 362 

alum, 302 

arsenite, 274 

aurate, 143 

bromide, 88 

carbonate, 331 

chlorate, 125 

chloride, 88 

chlorocbromate, 183 

chloroplatinate, 430 

chromate, 181 

chromic sulphate, 179 

cobalticyanide, 417 

cyanate, 344 

cyanide, 343 

diantimonate, 280 

dichromate, 182 

disulphate, 174 

ferricyanide, 412 

ferrocj^anide, 411 

fluoride, 88 

hydrate, 139 

hydroxide, 139 

hydrosulphide, 152 

iodide, 88 

isocyanate, 344 

manganate, 129 

nitrate, 236 

nitrite, 241 

oxides, 139 

per chlorate, 127 



INDEX. 



457 



Potassium permanganate, 130 

platinonitrite, 430 

silicates, 369 

stannate, 383 

sulphate, 173 

sulphides, 152 

sulphites, 158 

tri-iodide, 88 
Powder of algaroth, 277 
Practical units, 4 
Precipitate, 65 
Pressure, 3 
Prussian blue, 413 
Prussiate of potash, 411, 412 
Prussic acid, 342 
Purple of Cassius, 104 
Pyrolusite, 129 
Pyrophosphoric acid, 268 
Quartz, 366 
Quicksilver, 214 
Radicals, 162 
Rain water, 115 
Realgar, 275 
Red lead, 389 
Rochelle salt, 331 
Rock crystal, 366 
Rhodium and compounds of, 422 
Rubidium, 89" 

chloride, 89 

hydroxide, 140 

sulphate, 174 
Ruthenium and compounds of, 

421 
Sal- Ammoniac, 228 
Saleratus, 332 
Sal-soda, 328 
Salt-cake, 172 
Salt, common, 86 
Saltpetre, 236 
Salts, 162 

acid, basic and normal, 165 

formation of, 166 
Samarium and compounds of, 287 
Sand, sandstone, 366 
Scandium and compounds of, 308 
Scheele's green, 274 
Sea water, 115 

silver in, 98 
Second group, 195 
Selenic acid, 171 
Selenious acid, 171 
Selenium, 147 



Selenium dioxide, 170 
Serpentine, 206 
Seventh group, 56 

constitution of oxygen com- 
pounds of, 131 

oxides and hydroxides of, 123 

summary of, 135 

valence of, 131, 135 
Siderite, 394 

Silica and varieties of, 366 
Silicates, 368 
Silicic acids, 367, 368 
Silicon, 363 

chloroform, 366 

dioxide. 366 

hydroxides, 367 

tetrachloride, 365 

tetrarluoride, 364 

tetrahydride, 364 

trichloride, 366 
Silver, 97 

alloys of, 99 

bromide, 100 

chloride, 100 

coins, 99 

cyanide, 343 

dioxide, 142 

fluoride, 100 

hydroxide, 142 

iodide, 100 

nitrate, 237 

nitrite, 241 

orthophosphate, 265 

oxide, 142 

plating, 99 

sulphate, 175 

sulphide, 153 

tetrantoxide, 142 
Silvering, 99 
Sixth group, 107 

summary of, 192 
Slaked lime, 199 
Smithsonite, 209 
Soda-ash, 328 
Soda bicarbonate, 329 

caustic, 138 
Sodium, 85 

acetate, 362 

aluminates, 300 

antimonite, 279 

aurous sulphide, 154 

borates, 295, 296 



458 



IKDEX. 



Sodium bromide, 87 

carbonate, 328 

chloraurate, 106 

chloride, 86 

chloroplatinate, 430 

diantimouate, 280 

disulphate, 173 

fluoride, 87 

hydrate, 138 

hydroxide, 138 

hyposulphite, 170 

iodide, 87 ■ 

metastannate, 384 

nitrate, 237 

oxides, 137 

phosphates, 265 

potassium tartrate, 331 

pyrophosphate, 268 

pvrcsulphate, 173 

silicates, 369 

stannate, 383 

sulphantimonate, 281 

sulpharseuate, 275 

sulphate, 172 

sulphides, 153 

sulphite, 158 

thiocarbonate, 339 

thiosulphate, 170 

zirconate, 375 
Solder, plumber's, 387 
Solids, 5 

Soluble ferric hydroxide, 407 
Solution, 115 
Spathic iron, 403 
Specific heat, 28 

determination of, 28, 30 
Specific gravities, 3 
Spectral analysis, 90 
Spectroscope, 92 
Spelter, 209 
Spirits of wine, 357 
Spring water, 115 
Stannates, 383 
Stannic acid, 383 

chloride, 381 

hydroxides, 383 

oxide, 382 

sulphide, 384 
Stannous chloride, 380 

oxide, 380 

nitrate, 381 

sulphate, 381 



Stannous sulphide, 381 
Steel, manufacture of, 397 

properties of, 399 
Stibine, 277 
Stibnite, 280 
Stone ware, 371 
Strontianite, 202, 334 
Strontium, 202 

carbonate, 334 

chloride, 203 

dioxide, 203 

hydroxide, 203 

nitrate, 239 

oxide, 203 

sulphate, 203 
Sugar of lead, 390 
Sulphocyanic acid, 345 
Sulphur, 144 

amorphorus insoluble, 146 

amorphous sol able, 145 

dichloride, 151 

dioxide, 154 

heptoxide, 154 

monochloride, 151 

monoclinic, 145 

rhombic, 144 

sesquioxide, 154 

trioxide, 158 

tetrachloride, 152 
Sulphuretted hydrogen, 148 
Sulphuric acid, 159 

constitution of, 167 

di-, 169 

monometa-, dimeta-, and or- 
tho-, 168 
Sulphuric anhydride, 158 

dichloride, 167 

hj-dro chloride, 166 
Sulphurous acid, 157 

anhydride, 154 

oxide, 154 
Sulphuryl chloride, 167 

hydroxychloride, 166 
Superphosphate of lime, 267 
Symbols, 47 
Synthesis, 64 
Talc, 206 

Tantalum and compounds of, 288 
Telluric acid, 172 
Tellurium, 147 

dioxide, 171 

tetrachloride, 152 



INDEX. 



459 



Tellurium trioxide, 171 
Tellurous acid, 171 
Temperature, 24 
Tetrathionic acid, 170 
Tetragonal system, 17 
Tetramethylammonium iodide, 
354 

hydroxide, 354 
Thallic compounds, 307 
Thallium, 306 
Thallous compounds, 306 
Thermometers, 25 

calibration of, 26 
Thiocarbonic acid. 339 
Thiocyanic acid, 345 
Thiosulphuric acid, 169 
Third group, 291 

summary of, 311 
Thorite, 376 

Thorium and compounds of, 376 
Tin, 378 

dichloride, 380 

dioxide, 382 

salt, 380 

stone, 382 

tetrachloride, 381 
Titanium and compounds of, 372 
Triclinic system, 22 
Trihydrocyanic acid, 344 
Trimethylamine, 354 
Trithionic acid, 170 
Tungsten and compounds of, 187 
Tungsten sodium bronze, 190 
Turnbull's blue, 413 
Type metal, 276 
Unit of heat, 27 

length, 2 

mass, 2 

time, 2 
Units derived, 2 

fundamental, 2 

practical, 4 
Uranium and compounds of, 190 
Urea, 340 
Valence, 80, 131 



Valence of seventh group, 131, 135 

periodicity in, 446 
Vanadium and compounds of, 283 
Vaporization, 29 
Vapors, 6 
Verdigris, 362 
Vinegar, 361 

Victor Meyer's apparatus, 42 
Water, 113 

action of chlorine on, 65 

analysis of, 117 

carbonated, 325 

composition of, 115 

decomposition by sodium, 51 

distilled, rain, sea, and 
spring, 115 

table of pressure of vapor of, 
113 

vapor, weight of saturated, 250 

synthesis of, 117 
Weight, 4 
White lead, 391 
White vitriol, 212 
Witherite, 334 
Wood spirit, 348 
Wood's fusible metal, 281 
Wrought iron, 395, 399 

manufacture of, 397 
Ytterbium and compounds of, 

311 
Yttrium and compounds of, 309 
Zinc, 209 

blende, 209, 212 

chloride, 210 

chloride, use in soldering, 211 

dust, 210 

hydroxide, 212 

nitrate, 239 

oxide, 212 

oxychlorides, 211 

sulphate, 212 

sulphide, 212 

vitriol, 212 
Zircon, 374 
Zirconium and compounds of, 374 



'465 



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